Chapter 11Thermochemistry

1. Let’s begin by previewing the chapter (Page 292).

2. We will partner read Pages 293-94

The Flow of energy - heat

Thermochemistry – concerned with the

heat changes that occur during chemical

reactions.

Energy – the capacity to do work or

supplying heat.

Chemical Potential Energy

the energy stored within the structural

units of chemical substances.

◦ Energy stored in bonds that is released when

bonds are broken during chemical change

Heat

◦ Represented by the letter q

◦ A type of energy that transfers from one

object to another because of a temperature

difference between them.

Endothermic vs. Exothermic

System – the part of the universe on

which you focus your attention.

Surroundings – the part of the universe

that includes everything else in the

universe.

Universe = system + surroundings

◦ Our focus is on how chemical reactions affect

their surroundings.

Law of Conservation of Energy

During any chemical or physical process,

energy is neither created nor destroyed.

All energy involved in a process must be

accounted for as work, stored energy, or

heat.

Endothermic Process

A process that absorbs heat from the

surroundings.

The system gains heat from the

surroundings, the thermometer will show

a DECREASE in temperature in the

surroundings.

THE THERMOMETER IS IN THE

SURROUNDINGS, WE STUDY THE

SYSTEM, THE CHEMICAL BONDS.

Exothermic Process

A process that releases heat to its

surroundings.

The system loses heat to the

surroundings, the thermometer will RISE!

Heat Capacity and Specific Heat

A calorie is the quantity of heat needed

to raise the temperature of 1 gram of

pure water by 1C.

A Joule is the SI unit of heat and energy,

named after the English physicist James

Prescott Joule.

One calorie = 4.184 Joules

Heat capacity

The amount of heat needed to increase

the temperatur of an object exactly by

one degree Celsius.

A cup of water would have a much bigger

heat capacity than a drop of water.

Specific Heat

The amount of heat needed to raise the

temperature of one gram of the

substance by one degree Celsius.

Water has a very high specific heat value:

1.00 cal/g-C or 4.184 J/g-C.

Calculating heat, q

We use this formula to calculate the

amount of heat energy exchanged

between the system and its surroundings: q = (mass in grams)x (specific heat) x (change in temperature)

q = m c DT

Let’s try the example problems

together.

Calorimetry

The accurate and precise measurement of

heat change for a chemical and physical

process.

It uses an instrument known as a

calorimeter.

Coffee – Cup Calorimeter for

Constant Pressure Processes

q = m c DT

Bomb Calorimeter for Constant

Volume Combustions

Enthalpy

The heat content of a system, denoted

with the symbol H.

When pressure is held constant, as in the

case of our experiments in the lab, the

Enthalpy of a reaction is equal to the heat.

◦ q = H at constant pressure

Exothermic Reactions: DH = Negative

Endothermic Reactions: DH = Positive

Relationship

DH = q = m c DT is used to solve

calorimetry problems.

Let’s try some example problems.

Thermochemical Equations

An equation that includes the heat change

associated with a chemical process.

Let’s try some examples together.

Enthalpy diagrams

Show the exothermic or endothermic

process as a function of heat content

between reactants and products.

Let’s draw some together.

Heat in Changes of State

Phase diagrams display the state of a

substance at various pressures and

temperatures and the places where

equilibria exist between phases.

Phase Diagrams

Phase diagrams display the state of a substance at

various pressures and temperatures and the

places where equilibria exist between phases.

Phase Diagrams

The AB line is the liquid-vapor interface.

It starts at the triple point (A), the point at which

all three states are in equilibrium.

Phase Diagrams

It ends at the critical point (B); above this critical

temperature and critical pressure the liquid and

vapor are indistinguishable from each other.

Phase Diagrams

Each point along this line is the boiling point of

the substance at that pressure.

Phase Diagrams

The AD line is the interface between liquid and

solid.

The melting point at each pressure can be found

along this line.

Phase Diagrams

Below A the substance cannot exist in the liquidstate.

Along the AC line the solid and gas phases are in equilibrium; the sublimation point at each pressure is along this line.

Phase Diagram of Water

Note the high critical

temperature and critical

pressure:

◦ These are due to the strong

van der Waals forces between

water molecules.

Phase Diagram of Water

The slope of the solid–

liquid line is negative.

◦ This means that as the

pressure is increased at a

temperature just below the

melting point, water goes

from a solid to a liquid.

Phase Diagram of Carbon Dioxide

Carbon dioxide

cannot exist in the

liquid state at

pressures below 5.11

atm; CO2 sublimes at

normal pressures.

Phase Changes

Energy Changes Associated with

Changes of State

Heat of Fusion: Energy required to change a

solid at its melting point to a liquid. DHfus

Energy Changes Associated with

Changes of State

Heat of Vaporization: Energy required to change a liquid at its boiling point to a gas. DHvap

Notice that the heat of vaporization is always larger than its heat of fusion.

Water

The heat of fusion, or enthalpy of fusion,

for ice is 6.01 kJ/mol.

The heat of vaporization, or enthalpy of

vaporization, for water is 40.7 kJ/mol.

The heat of sublimation is the sum of

heats of vaporization and fusion.

◦ For water = approx 47 kJ/mol

Energy Changes Associated with

Changes of State

The heat added to the

system at the melting and

boiling points goes into

pulling the molecules farther

apart from each other.

The temperature of the

substance does not rise

during the phase change.

Calculating DH for Temperature and

Phase Changes

Calculate the enthalpy change upon converting 1 mol of ice at -25oC to water vapor (steam) at 125oC under a constant pressure of 1 atm.

◦ The specific heats of ice, water, and steam are 2.09 J/g-K, 4.18 J/g-K, and 1.84 J/g-K respectively.

◦ For H2O, DHfus = 6.01 kJ/mol and DHvap = 40.67 kJ/mol.

Practice Exercise

What is the enthalpy change during the

process in which 100 g of water at

50.0oC is cooled to ice at -30oC?

Heats of Solution

Heat changes can also occur when a

solute dissolves in a solvent.

The heat change caused by dissolution of

one mole of substance is the Molar Heat

of Solution, DHsol.

Let’s write some together, and work the

practice problems.

Hess’s Law of Heat Summation

If you add two or more thermochemical

equations to give a final equation, then

you can also add the heats of reaction to

give the final heat of reaction.

Two rules

◦ If the reaction is reversed the sign of DH is

changed

◦ If the reaction is multiplied or divided, so is

DH

Using Hess’s Law to

Calculate DH The following information is known:

◦ C(s) + O2(g) CO2(g) DH1 = -393.5 kJ

◦ CO(g) + ½ O2 (g) CO2 (g) DH2 = -283.0

kJ

Using these data, calculate the enthalpy

for:

◦ C(s) + ½ O2(g) CO(g)

More Practice with Hess’s Law

Calculate DH for the reaction

◦ 2C(s) + H2(g) C2H2(g)

Given the following chemical equations

and their respective DH.

◦ C2H2(g) + 5/2O2 2CO2(g) + H2O (l)

DH = - 1299.6

kJ

◦ C(s) + O2(g) CO2(g) DH = -393.5 kJ

◦ H2(g) + ½ O2(g) H2O(l) DH = -285.8 kJ

You Try It…

Calculate DH for the reaction

◦ NO(g) + O(g) NO2 (g)

Given the following information:

NO(g) +O3(g) NO2(g) + O2(g) DH = -198.9kJ

O3(g) 3/2 O2(g) DH = -142.3 kJ

O2(g) 2 O (g) DH = 495.8 kJ

Remember…

H is a state function, so for a particular

set of reactants and products, DH is the

same whether the reaction takes place in

one step or in a series of steps.