William L MastertonCecile N. Hurleyhttp://academic.cengage.com/chemistry/masterton

Edward J. Neth • University of Connecticut

Chapter 8Thermochemistry

Heat (Not in Notes)

• Heat will flow from a hotter object to a colder object• Mix boiling water with ice• Temperature of the ice rises after it melts• Temperature of the water falls

8.1 Principles of Heat Flow

Thermodynamics: the study of energy changes in chemical reactions and the influence of energy on those changes.

system: that part of the universe on which attention is focused

surroundings: the rest of the universe

Heat: thermal energy transferred between the system and the surroundings

Temperature: a measure of the average kinetic energy of the particles in a system

Figure 6.1 – Systems and Surroundings

Chemical Reactions (Not in Notes)

• When we study a chemical reaction, we consider the system to be the reactants and products

• The surroundings are the vessel (beaker, test tube, flask) in which the reaction takes place plus the air or other material in thermal contact with the reaction system

I. Direction of Heat Flow: Heat is given the symbol, q

• q is positive when heat flows into the system from the surroundings

• q is negative when heat flows out of the system into the surroundings

• Endothermic processes have positive q

• H2O (s) H2O (ℓ) q > 0

• Exothermic processes have negative q

• CH4 (g) + 2O2 (g) CO2 (g) + H2O (ℓ) q < 0

Exothermic and Endothermic Processes

II. Magnitude of Heat Flow (q) – the amount of heat transfer

James Joule: (1818-1889) calorimetry

calorie: English unit of measure for heat; 1 calorie is equal to the amount of energy required to increase the temperature of a 1g sample of water 1°C

Calorie: Nutritional Calorie; 1,000cal = 1Cal

Joule: S.I. unit – Joules(J) or kilojoules(kJ) are the metric measures of heat

Conversions (calories to Joules):• 1 calorie = 4.184 J• 1 kilocalorie = 1 Calorie ( nutritional calorie)• Nutritional calories are kcal

III. The Calorimetry Equation

• q = C x t

• t = tfinal – tinitial

• C (uppercase) is the heat capacity of the system: it is the quantity of heat needed to raise the temperature of the system by 1 °C

• q = m x c x t• c (lowercase) is the specific heat: the quantity of

heat needed to raise the temperature of one gram of a substance by 1 °C

• c depends on the identity and phase of the substance

Specific Heat

• The specific heat of a substance, like the density or melting point, is an intensive property that can be used to identify a substance or determine its purity

• Water• Water has an unusually large specific heat• A large quantity of heat is required to raise the

temperature of water• Climate is moderated by the specific heat of water

Table 8.1

Example 8.1

8.2 Measurement of Heat Flow

• Calorimeter: a device used to measure the heat flow of a reaction• The walls of the calorimeter are insulated to block

heat flow between the reaction and the surroundings

• The heat flow for the system is equal in magnitude and opposite in sign from the heat flow of the calorimeter• qreaction = - qcalorimeter

• qreaction = - Ccal Δt

Figure 8.2

Coffee-cup Calorimeter

• For a reaction performed in a coffee-cup calorimeter

tCg

Jmq waterreaction

18.4

Example 8.2

Heat calculation for mixing of solutions:

• If the concentrations are less than or equal to 1.0M, you can assume the density to be the same as water 1g/mL, and you must add the masses together to get the total mass of the combined solution.

Figure 8.3

Bomb Calorimeter

• The bomb calorimeter is more versatile than the coffee-cup calorimeter• Reactions involving high temperatures and gases

• A heavy metal vessel surrounded by water

• qreaction = -qcalorimeter

• qreaction = -CcalΔt

• Ccal is specific to each calorimeter and can be measured experimentally

Enthalpy

• The heat flow at constant pressure is equal to the difference in enthalpy (heat content) between products and reactants

• The symbol for enthalpy is H• We measure changes in enthalpy using a

calorimeter and a reaction run at constant pressure:• ΔH = Hproducts – Hreactants

• The sign of the enthalpy change is the same as for heat flow:• ΔH > 0 for endothermic reactions• ΔH < 0 for exothermic reactions• Enthalpy is a state variable

Exothermic Reactions

Figure 8.4 – Enthalpy of Reaction

Thermochemical Equations• A thermochemical equation is a chemical equation with the ΔH for the reaction included

• Endothermic Example – energy is added to the system• NH4NO3 (s) NH4

+ (aq) + NO3- (aq) H = +28.1 kJ

Exothermic Example – energy is released to the system

CH4(g) +2O2(g)CO2(g) + 2H2O(l) + 891kJ

Thermochemical equation

CH4(g) +2O2(g)CO2(g) + 2H2O(l) ∆H= -891kJ

Figure 8.5 – An Endothermic Reaction

Conventions for Thermochemical Equations

1. The sign of H indicates whether the reaction is endothermic or exothermic

2. The coefficients of the thermochemical equation represent the number of moles of reactant and product

3. The phases of all reactant and product species must be stated

4. The value of H applies when products and reactants are at the same temperature, usually 25 °C

Rules of Thermochemistry

1. The magnitude of H is directly proportional to the amount of reactant or product

2. H for the reaction is equal in magnitude but opposite in sign for H for the reverse of the reaction

3. The value of H is the same whether the reaction occurs in one step or as a series of steps

This rule is a direct consequence of the fact that ΔH is a state variable

This rule is a statement of Hess’s Law

Example 8.4

Enthalpy of Phase Changes

• Phase changes involve enthalpy• There is no change in temperature during a phase change• Endothermic: melting or vaporization• Exothermic: freezing or condensation

• Pure substances have a value of ΔH that corresponds to melting (reverse, fusion) or vaporization (reverse, condensation)

Hess’s Law:

• The heat of a reaction (H ) is constant, whether the reaction is carried out directly in one step or indirectly through a number of steps.

• The heat of reaction (H ) can be determined as the sum of the heats of reaction of several steps

Example 8.5

Example 8.6

8.5 Enthalpies of Formation

• The standard molar enthalpy of formation, , is equal to the enthalpy change• For one mole of a compound • At constant pressure of 1 atm • At a fixed temperature of 25 °C • From elements in their stable states at that temperature

and pressure• Enthalpies of formation are tabulated in Table 8.3 and in

Appendix 1 in the back of the textbook

fH

fH

• The standard enthalpy of formation of a pure element in its standard state at 25 °C is zero

• The enthalpy of formation of H+ (aq) is also zero

Calculation of

• The symbol Σ refers to “the sum of”

• Elements in their standard states may be omitted, as their enthalpies of formation are zero (standard state only)

• The coefficients of reactants and products in the balanced equation must be accounted for

reactantsHproductsHH ff

H

Example 8.7

Example 8.8

8.7 The First Law of Thermodynamics

• Thermodynamics• Deals with all kinds of energy effects in all kinds of

processes• Two types of energy• Heat (q) • Work (w)

• The Law of Conservation of Energy• E system = - Esurroundings

• The First Law• E = q + w• The total change in energy is equal to the sum of the

heat and work transferred between the system and the surroundings

Conventions

• q and w are positive• When the heat or work enters the system from the

surroundings• q and w are negative • When the heat or work leaves the system for the

surroundings

Figure 8.10

Example 8.9

Chapter 17 Spontaneity of Reactions

• 17.1 Spontaneous Processes - Everyday process that take place on their own, without outside forces

• What is a spontaneous process? (examples)• An ice cube will melt when added to a glass of

water at room temperature• A mixture of hydrogen and oxygen will form water

when a spark is applied• An iron (or steel) tool will rust if exposed to moist

air

Spontaneity

• Spontaneity and rate are not connected; FAST DOES NOT = SPONTANEOUS

• If a reaction is spontaneous in one direction, it will be non-spontaneous in the reverse direction under the same conditions

Spontaneity and Equilibrium

• A spontaneous process is one that moves a reaction system toward equilibrium, and a nonspontaneous process moves away from equilibrium

(equilibrium = a state of no net change in a process)

Which of the following are spontaneous processes?

1. Snowman melting in the sun

2. Assembling a jigsaw puzzle

3. Rusting of an iron object in humid air

4. Recharging of a camera battery

The Energy Factor

• Many spontaneous reactions are exothermic, but not all!

• Many spontaneous processes proceed with a decrease in energy• Boulders roll downhill• Your cell phone battery discharges over time

• Recall that exothermic reactions proceed with a decrease in energy• Spontaneous reactions are typically exothermic

Exceptions

• phase changes

• H2O (s) H2O (l) is endothermic but spontaneous at room temperature

• Some reactions become spontaneous with a simple increase in temperature• CaCO3 (s) CaO (s) + CO2 (g) ΔH = +178.3 kJ

• ΔH is not the only criterion for spontaneity

The Randomness Factor

• Nature tends to move spontaneously from a state of lower probability (order) to one of higher probability (disorder), or

• Systems overtime without outside influence, will move toward a condition of maximum probability (disorder).

Roll of the Dice (NOT IN NOTES – FOR DISCUSSION)

• When rolling a pair of dice:• There is only one way to roll a 2 or a 12• There are six ways to roll a 7• The probability of rolling a 7 is six times greater

than that of rolling 2 or 12• The state 7 is of higher probability than the state 2

or 12

Figure 17.2

17.2 Entropy, S

What is entropy?• described as an increase in disorder or randomness;

symbol or variable S.• A measure of the amount of randomess in a

system• When water vaporizes the entropy of the gas

molecules is greater than the liquid• Kinetic molecular theory supports greater motion

will result in greater entropy

Figure 17.3 – Disorder and Order

Factors that Influence the amount of entropy

• A liquid has higher entropy than the solid from which it formed

• A gas has higher entropy than the liquid or solid from which it formed

• Increasing the temperature of a substance increases entropy

Figure 17.4

The Third Law of Thermodynamics

• A completely ordered, pure crystalline solid at 0 K has an entropy of zero• This is the only time the entropy of a substance is

zero• Absolute zero has not been reached; it is still a

theoretical limit

Example 17.1

Standard Molar Entropies

• Unlike enthalpy, molar entropy cannot be directly measured• Elements have nonzero molar entropies• In calculating the standard molar entropy change,

elements must be taken into account

• Standard molar entropies are always positive numbers, i.e., S° > 0

• Aqueous ions may have negative S°values

ΔS° for Reactions

• Units are J/mol-K or kJ/mol-K

reactants,S,SS products

The Second Law of Thermodynamics

• In a spontaneous process, there is a net increase in entropy, taking into consideration both the system and the surroundings

• That is, for a spontaneous process:

0)( gssurroundinsystemuniverse SSS

17.3 Free Energy, G

• Gibbs Free Energy• If G is negative,• < 0, the reaction is spontaneous

• If G is positive,• > 0, the reaction is nonspontaneous as written (the

reverse reaction is spontaneous)

• If G = 0, • the system is in equilibrium; no tendency for the

reaction to occur in either direction

Figure 17.5

Relation Among ΔG, ΔH and ΔS

• The Gibbs-Helmholtz equation• G = H – T S

• Spontaneous reactions generally have• H < 0 and S > 0

• In specific cases, either term may dominate• phase changes, S is dominant; some reactions,

S is nearly zero and H will dominate

17.4 The Standard Free Energy Change, ΔG°

• Standard conditions:• Gases are at 1 atm pressure• Solutions are 1M for ions or molecules

• Under standard conditions:• G° = H° – T S°

• Recall that• If ΔG° < 0, spontaneous• If G° > 0, nonspontaneous• If G° = 0, equilibrium

Free Energy of Formation

• We can use the Gibbs-Helmholtz equation to calculate the standard free energy of formation for a compound

The sign of • If negative, the formation of the compound is spontaneous• If positive, the formation of the compound is nonspontaneous

fG

Example 17.3

17.5 Effect of Temperature, pressure and concentration on reaction spontaneity.

Table in notes

Table 17.2

Example 17.4

Predict which of the four cases in the table above will apply to the following reactions, and if the reaction will be spontaneous. Use ∆H as given and your estimate of the sign of ∆S.

C6H12O6 (s)+6O2(g)6CO2(g)+6H2O(g)

ΔH = -2540 kJ

Cl2(g) 2Cl(g)

ΔH is positive

17.5 Standard Free Energy Change

• A. Equation at standard conditions:

ΔGo = ΔHo - T ΔSo