© 2009, Prentice-Hall, Inc. Chapter 17 Thermochemistry Edited by Kavita Gupta Chemistry, The Central Science, 11th edition Theodore L. Brown; H. Eugene

© 2009, Prentice-Hall, Inc.

Chapter 17Thermochemistry

Edited by Kavita Gupta

Chemistry, The Central Science, 11th editionTheodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten


Energy• Energy is the ability to do work or transfer

heat.– Energy used to cause an object that has mass to

move is called work.– Energy used to cause the temperature of an object

to rise is called heat.– There are different forms of energies- radiant,

thermal, chemical, potential etc. We will focus on thermal (heat) energy in this unit.

• Units for energy: • The SI unit of energy is the joule. Another

common unit of energy is calorie. • 1 calorie = 4.184 joules


Heat

• Energy can also be transferred as heat.

• Heat flows from warmer objects to cooler objects.

Temperature

TEMPERATURE: a measure of average kinetic energy of the particles of matter (remember that molecules of ALL gases at the same temp. have the same average kinetic energy)

THERMOCHEMISTRY: Branch of Chemistry that deals with the study of heat change in chemical reactions.


Definitions:System and Surroundings

• The system includes the molecules we want to study (here, the hydrogen and oxygen molecules).

• The surroundings are everything else (here, the cylinder and piston).

System

SYSTEM: that part of the universe on which attention is focused. Usually reactants and products.

There are three types of systems:• Open System: can exchange both mass and energy.• Closed System: can exchange only energy.• Isolated System: can neither exchange mass nor

energy. • What kind of system is a heat/cold pack?

Surroundings

• SURROUNDINGS: what exchanges energy with the system, the vessel in which the reaction takes place and the air or other material in contact with the reaction system.

• Label system and surrounding in the following diagram.


Exchange of Heat between System and Surroundings

• When heat is absorbed by the system from the surroundings, the process is endothermic. (usually the temp of the system increases and it feels cold to touch!)


Exchange of Heat between System and Surroundings

• When heat is absorbed by the system from the surroundings, the process is endothermic.

• When heat is released by the system into the surroundings, the process is exothermic. (Usually the temp. of the system lowers and it feels warm to touch)


State Functions

Usually we have no way of knowing the internal energy of a system; finding that value is simply too complex a problem.


Enthalpy• Enthalpy (H) is defined as the heat content of a system at constant

pressure. • It is an extensive property.• It is not possible to measure the heat content of the system only the

changes in the heat content of the system can be measured. Enthalpy change is defined as the amount of heat absorbed or lost by a system during a process at constant pressure.

• Heat of reaction and change in enthalpy are used interchangeably. • Units: kJ/mol• Qreaction at constant pressure= H(change in enthalpy) = H (products)

– H (reactants)


Endothermicity and Exothermicity

• A process is endothermic when H is positive.


Endothermicity and Exothermicity

• A process is endothermic when H is positive.

• A process is exothermic when H is negative.

Energy Diagram for Endo- and Exothermic Reactions

• Endothermic (dH +) • Exothermic ( dH -)


Enthalpy of Reaction

The change in enthalpy, H, is the enthalpy of the products minus the enthalpy of the reactants:

H = Hproducts − Hreactants


Enthalpy of Reaction

This quantity, H, is called the enthalpy of reaction, or the heat of reaction.


The Truth about Enthalpy1. Enthalpy is an extensive property. hence it is

proportional to the amount of reactants and products. e.g. for decomposition of two moles of water twice as much energy is needed as for one mole of water.

2. H for a reaction in the forward direction is equal in size, but opposite in sign, to H for the reverse reaction. Reversing a reaction changes the sign of enthalpy. AB 10kJ, then BA –10kJ

3. H for a reaction depends on the state of the products and the state of the reactants. (different values of dH for H2O (l) and H2O (g).

Thermochemical Equation

• Thermochemical equation: A chemical equation that shows the enthalpy relation between products and reactants is called a thermochemical equation.

• Ex. H2 (g) + Cl2 (g)2 HCl (g) H=-185 kJ


Heat Capacity and Specific Heat

The amount of energy required to raise the temperature of a substance by 1 K (1C) is its heat capacity.



We define specific heat capacity (or simply specific heat) as the amount of energy required to raise the temperature of 1 g of a substance by 1 K.



Heat Capacity (C)

Formula for Specific Heat q(amount of heat)= m (g of substance) x c(specific heat) x T ( temperature change)


Calorimetry

Since we cannot know the exact enthalpy of the reactants and products, we measure H through calorimetry, the measurement of heat flow.

Calorimeter

• Calorimeter is used to measure heat flow in a reaction.. Only heat flow is between the reaction system and the calorimeter, the walls of the calorimeter is insulated, so no heat exchange occurs between calorimeter and surroundings.

• q reaction = - q calormeter

• if reaction is exothermic, heat flows into calorimeter, if reaction is endothermic heat flows into reaction mixture..

Types of Calorimeters• There are two types of calorimeters:• Coffee-Cup Calorimeter (Constant Pressure

Calorimetry): It consists of foam cup filled with water. Cup has a tight fitting lid through which a thermometer is inserted. Because Styrofoam is a good insulator, very little heat is lost to the surroundings. This means that the heat capacity of calorimeter is that of the water, which can be calculated by

Qrxn=-mwater x 4.18 j/g-0c x tCoffee cup calorimeter can be used to measure

heat of reaction of solutions only


Constant Pressure Calorimetry

By carrying out a reaction in aqueous solution in a simple calorimeter such as this one, one can indirectly measure the heat change for the system by measuring the heat change for the water in the calorimeter.


Constant Pressure Calorimetry

Because the specific heat for water is well known (4.184 J/g-K), we can measure H for the reaction with this equation:q = m s T

Bomb Calorimeter

• Bomb Calorimeter: Bomb calorimeter is used to measure heat of reactions for gases and also for reactions that reach high temperatures. To use it, a weighed sample of a reactant is put in a heavy metal vessel called as a “bomb”. Initial temperature is recorded and reaction is started by an ignition. Final temperature is recorded.

• Qrxn= -q calorimeter


Bomb Calorimetry

• Reactions can be carried out in a sealed “bomb” such as this one.

• The heat absorbed (or released) by the water is a very good approximation of the enthalpy change for the reaction.


Bomb Calorimetry

• Because the volume in the bomb calorimeter is constant, what is measured is really the change in internal energy, E, not H.

• For most reactions, the difference is very small.

Heat Transfer (Heating and Cooling Curves)

• Heat Transfer :Look at the following heating/cooling curve for water (Remember this from Chap 12)

• Label on this graph for: temperatures, fusion, freezing,

vaporization, condensation• http://www.wwnorton.com/college/chemistry/gilbert/tuto

rials/interface.asp?chapter=chapter_11&folder=heating_curves (Heating/Cooling Curve Animation)

http://www.wwnorton.com/college/chemistry/gilbert/tutorials/interface.asp?chapter=chapter_11&folder=heating_curves



For Calculating Energy at Phase Change

• During phase changes all the heat is used for

breaking the bonds. To calculate amount of heat absorbed/released during phase changes (q), the following formula is used: Q = m. H ( or n. H)

• Value of H changes with phase change. H of water has following values:

• Hfusion = 334 J/g (6.02 kJ/mol), Hfreezing = - 334 J/g (-6.02 kJ/mol),

• Hvaporization = 2260 J/g (40.7 kJ/mol), Hcondensation = -2260 J/g (-40.7 kJ/mol)

For Calculating Energy other than at phase change

• Use the formula Q=m x c x t

• Total energy change can be calculated by adding all the energy changes.

Ex. How much heat is needed to vaporize 5.0 g of ice at -5 degree C to steam at 110 degree C?

• Enthalpy of Formation(Hf), of a substance is the enthalpy change for the reaction in which the substance is formed from its constituent elements.

Write an equation for heat of formation of Carbon dioxide gas.

Write a thermochemical equation for the formation of AgCl, if Hf

0 for AgCl is –127.1 kJ/mol.

• Standard enthalpy of formation, Hf0,of a substance is

the change in enthalpy for the reaction that forms 1 mole of a substance from its elements with all the reactants and products1 atm pressure and 298 K.

• For a pure element in its most stable

state,Hf0=0• Compounds with high negative enthalpy of

formation are very stable while the ones with high positive enthalpy of formation are unstable. Why?

• The importance of standard enthalpies of formation is that once we know their values, we can calculate standard enthalpy of reaction, H0. How?


Hess’s Law

• H is well known for many reactions, and it is inconvenient to measure H for every reaction in which we are interested.

• However, we can estimate H using published H values and the properties of enthalpy.


Hess’s Law

Hess’s law states that “[i]f a reaction is carried out in a series of steps, H for the overall reaction will be equal to the sum of the enthalpy changes for the individual steps.”


Hess’s Law

Because H is a state function, the total enthalpy change depends only on the initial state of the reactants and the final state of the products.


Enthalpies of Formation

An enthalpy of formation, Hf, is defined as the enthalpy change for the reaction in which a compound is made from its constituent elements in their elemental forms.


Standard Enthalpies of Formation

Standard enthalpies of formation, Hf°, are measured under standard conditions (25 °C and 1.00 atm pressure).


Calculation of H

We can use Hess’s law in this way:

H = nHf°products – mHf° reactants

where n and m are the stoichiometric coefficients.


H = [3(-393.5 kJ) + 4(-285.8 kJ)] – [1(-103.85 kJ) + 5(0 kJ)]= [(-1180.5 kJ) + (-1143.2 kJ)] – [(-103.85 kJ) + (0 kJ)]= (-2323.7 kJ) – (-103.85 kJ) = -2219.9 kJ

C3H8 (g) + 5 O2 (g) 3 CO2 (g) + 4 H2O (l)

Calculation of H

Hess’s Law Contd.• Many compounds can not be synthesized directly from their elements. In such

cases, Hess’s law is used to calculate H of the reaction.• Hess’s Law:• According to Hess’s law, if a reaction is carried out in a series of steps, H for

the reaction will be equal to the sum of the enthalpy changes for the steps. Example:Use Hess's law to calculate H for the reaction: MgO(s) + CO2(g) MgCO3(s) from the following enthalpy of formation data.Mg(s) + 1/2 O2(g)MgO(s) Hf

0 = -601.70 kJ/mol MgO C(s) + O2(g)CO2(g) Hf

0= -393.51 kJ/mol CO2 Mg(s) + C(s) + 3/2 O2(g)MgCO3(s) Hf

0 = -1095.8 kJ/mol MgCO3 • Do practice problems1,2,3 on page 524

Enthalpy

• Enthalpy of Combustion is the heat released by the complete combustion of one mole of a substance. Hc

• • Example: Write a thermochemical equation

for combustion of C6H6, if Hc for C6H6 is –3169.0 kJ/mol.

•

Enthalpy of Vaporization

• Enthalpy of Vaporization: Amount of heat required to boil (or condense) 1 mol of a substance at its boiling point. Hvap for water= 40.7 kJ/mol, what would Hcondensation be for water?

Example:• How much heat is required to vaporize 5

moles of water at its B.P.?

Enthalpy of Fusion

• Enthalpy of Fusion: Amount of heat required to melt (or freeze) 1.00 mol of a substance at its freezing point. Hfus for water=6.02kJ/mol

Spontaneity of A Chemical Reaction:• Spontaneity of a reaction refers to if the reaction occurs

spontaneously. Spontaneous reactions may not always be fast. Some examples of spontaneous reactions are rusting, reaction of baking soda with vinegar etc.

• Non-Spontaneous Reactions: need something to get them started like heat, pressure or catalyst etc.

• There are two factors that affect the spontaneity of a reaction:

• Enthalpy (H), exothermic reactions (-H) are usually spontaneous while endothermic reactions (+H) are usually non spontaneous. However, some endothermic reactions are spontaneous such as melting.

• Entropy: Reaction is spontaneous, if value of S is positive and non spontaneous, if value of S is negative.

Entropy (S) • It is a measure of randomness or disorder of the particles in a

system. Units: kJ/mol-K• General rules for predicting entropy changes. 1. Look at the states first. (gases > liquids > solids) 2. If both states are the same then look at the number of moles of

reactants and products and decide if there has been an increase in the number of moles or a decrease.

3. If the volume of a container of gas is increased, it results in increased entropy (more space for random movement)

4. If the pressure on a container of gas is increased, then the entropy decreases (due to restricted movement of particles)

• Arrange ice, water and water vapor in order of decreasing entropy.

Gibbs Free Energy

• Gibbs Free Energy(or Free Energy) G, is a

thermodynamic quantity that combines the values of enthalpy and entropy. For a process that occurs at constant temperature,

• G=H-TS

• The sign of G relates to the spontaneity of the reaction. When it is negative process is spontaneous.

Relating Enthalpy, Entropy, Gibbs free energy to the spontaneity of the reactions

H S G

Negative

value(exothermic)

Positive value(disorder)

Always negative

( spontaneous)

Negative value

Negative value Negative at lower

temperatures

Positive value

Positive value Negative at higher

temperatures

Positive value

Negative value Never negative (non

spontaneous)

Kinetics• Kinetics • What is Chemical Kinetics? • Chemical kinetics is the study of rates of reaction.• Rate: measure of the speed of any change that occurs within an

interval of time. • EX: A sprinter covers 100 meters in 11.5 seconds. Rate is 8.7 m/s. • 1. Reaction X O / min, then, 1 min. XXXXXO 2 min. XXXXOO 3

min. XXXOOO 4 min. XXOOOO 5 min. XOOOO

Reaction Rates in Chemistry • Reaction Rates in Chemistry are usually described in terms of change in

concentration per unit time. It can be written either as appearance of product per unit time or disappearance of reactants per unit time. Units: M/s. Reaction rates can be determined experimentally by changing concentration of reactants. For example, rate of reaction for the following reaction can be written in terms of disappearance of HI or appearance of H2 or I2.

• 2 HI (g) H2(g) + I2 (g)

• Also, note that the rate of change of HI is twice as compared to formation of H2 or I2. So, we can equate these rates as follows:

• Units for rate: change in concentration/ change in time (Molarity/second

or Ms-1)

Rate Laws• Rate Laws of Reactions:• Rate law is an equation that shows quantitative

relationship between the reaction rate and concentrations of reactants. This relationship can only be measured experimentally.

• r [A], r k[A], where k is specific rate constant. The

value of k changes only with temperature. • To calculate rate law, the concentration of reactants

are changed and changes in rate is observed experimentally.

Experiment # [A] [B] Rate (M/s)

1.100 .100 4.0 X 10^-5

2 .100 .200 8.0 X 10^-5

3 .200 .100 1.6

To calculate rate law, the concentration of reactants are changed and changes in rate is observed experimentally.

• Ex: 2H2(g) + 2NO(g) N2(g) + 2 H2O(g)

• In the above example, 4 moles of reactant produces 3 moles of product gases, pressure goes down as reaction proceeds. Rate determined by measuring change in pressure over time.

• In the above example, 4 moles of reactant produces 3 moles of product gases, pressure goes down as reaction proceeds. Rate determined by measuring change in pressure over time.

• 1. If start with same initial conc of NO but different H2 . Initial reaction rate is found to vary directly with the hydrogen conc. If double H conc doubles rate, if triple H2 triples rate, IF r= react rate and (H2) = conc of H2 in moles/L. then r is proportional, , to (H2). R (H2)x 2 (2)x r=1

• if use same concentration of H2 but vary NO . Initial reaction rate is

found to increase 4X when NO conc is doubled and nine fold when conc is tripled.. Reaction rate varies directly with the square of the nitrogen monoxide conc. R (NO)2. R (NO)x 4 (2)x r=2

• r (H2)(NO) 2

(To remove the proportionality sign, a constant called as specific rate constant (k) is introduced.)

k is proportionality constant. So r= k(H2)(NO) 2

k is specific for a given reaction at a given temp,

k usually increases with increase in temp.

Order of Reaction• Order of Reaction: • The power to which a reactant conc is raised is called the

ORDER in that reactant..• An order of 1, rate is directly proportional.• 2, rate is directly proportional to the square of

the reactant• 0, rate does not depend on the concentration

of the reactant as long as some of the reactant is present. • Overall order of reaction: sum of order of reactants. • Order with respect to H2 is first order, with respect to NO is

second order and to overall reaction is 3rd.

Units of Rate Constants• In zero order reaction,the rate of reaction is

independent of the concentration of reactants.• Rate= k units of k= M/s • In first order reaction, the rate is proportional to

the concentration of only one of the reactants raised to the first power.

• Rate= k[A] units of k= s-1• In second order reaction, the overall order of

reaction is 2,• Rate= k[A][B] or Rate= k [A]2 units: M-1s-1

Reaction Pathways for Forward and Reverse Reactions: An Energy Diagram

Forward reaction Reverse reaction exothermic endothermic

For Forward reaction: w =reactantsa =activation energy, Ea

y = activated complex c= energy change in the reaction, Delta Ez = productsb = energy needed to activate the reverse reaction ( endothermic) Ea

’

Reaction Mechanism: A reaction mechanism details the individual steps that occur in the course of a reaction.Each of these steps are called as elementary steps. An elementary step may produce an intermediate, a product that is consumed in a later elementary step and therefore does not appear in the overall stoichiometry of the reaction. If a mechanism has several elementary steps, then the overall rate is determined by t he slowest step, called the rate-determining step.

Collision Theory: Atoms, ions and molecules can react to form products when they collide with one another, if they have 1.enough KE and 2. the right orientation.

Activation Energy: Ea, The minimum amount of energy that particles must have in order to react.

Activated complex: an unstable arrangement of atoms that forms momentarily at the peak of the activation energy.

Factors Affecting Reaction Rate:1.Temperature: increase temp, increase collisions, & increase # collisions with enough kinetic energy to get over Ea barrier.2.Concentration: increase concentration increases frequency of collisions leads to higher reaction rate.3.Particle Size: increase in surface area increases amount of reactant exposed, increases collision frequency and reaction rate.4.Catalysts: increases rate of reaction without being used up during the reaction. Lower activation energy barrier so more reactants have the energy to form products within a given time.

Documents

© 2009, Prentice-Hall, Inc. Chapter 17 Thermochemistry Edited by Kavita Gupta Chemistry, The Central Science, 11th edition Theodore L. Brown; H. Eugene