Unit 6: Thermochemistry

Unit 6: Thermochemistry

Introduction Heat and Work Specific Heat Enthalpy ( Enthalpy of Reaction Phase Diagram

Introduction

Most daily activities involve processes that either use or produce energy:Activities that produce energy

Metabolism of foodBurning fossil fuels

Activities that use energy:PhotosynthesisPushing a bike up a hillBaking bread

Introduction

ThermodynamicsThe study of energy and its

transformations

Thermochemistry:A branch of thermodynamicsThe study of the energy (heat)

absorbed or released during chemical reactions

Introduction

Objects can have two types of energy:Kinetic energy

Energy of motionThermal energy

The type of kinetic energy a substance possesses because of its temperature

Potential energyEnergy of position“stored” energy resulting from the attractions and repulsions an object experiences relative to other objects

Introduction

Units of EnergySI unit = joule (J)

1 J = the kinetic energy of a 2 kg mass moving at a speed of 1 m/s

A very small quantity

Kilojoule (kJ)1 kJ = 1000 J

Introduction

Units of Energy (cont)Calorie (cal)

Originally defined as the amount of energy needed to raise the temperature of 1g of water from 14.5oC to 15.5oC.

1 cal = 4.184 J (exactly)

Kilocalorie (kcal)1 kcal = 1000 cal

Introduction

Example: Convert 3.02 kJ to J.

Given: 3.02 kJFind: J

J = 3.02 kJ x J = 3.02 kJ x 1000 J1000 J = 3020 J = 3020 J1 kJ1 kJ

1 kJ = 1000 J

Introduction

Example: Convert 725 cal to kJ.

Given: 725 calFind: kJ

J = 725 cal x J = 725 cal x 4.184 J4.184 J x x 1 k J1 k J = 3.03 kJ = 3.03 kJ1 cal1 cal 1000 J1000 J

1 cal = 4.184 J1 kJ = 1000 J

Introduction

When using thermodynamics to study energy changes, we generally focus on a limited, well-defined part of the universe.

System:The portion of the universe

singled out for study

Surroundings:Everything else

Introduction

The system

The system is usually the chemicals in the flask/reactor. The flask and everything else belong to the surroundings.

Introduction

Open system:A system that can exchange both

matter and energy with the surroundings

Closed system:A system that can exchange

energy with the surroundings but not matter

A cylinder with a piston is one example of a closed system.

Introduction In a closed system

energy can be gained from or lost to the surroundings as:

WorkHeat

Work:Energy used to cause

an object to move against a forceLifting an object Hitting a baseball

Introduction

Heat:The energy used to cause the

temperature of an object to increase

The energy transferred from a hotter object to a cooler one

Energy:The capacity to do work or to

transfer heat

Introduction

The potential energy of a system can be converted into kinetic energy and vice versa.

Energy can be transferred back and forth between the system and the surroundings as work and/or heat.

Potential energy Kinetic energy

work

The First Law of Thermodynamics

Although energy can be converted from one form to another and can be transferred between the system and the surroundings:

Energy cannot be created or destroyed.(First Law of Thermodynamics)

Any energy lost by the system must be gained by the surroundings and vice versa.


The First Law of Thermodynamics can be used to analyze changes in the Internal Energy (E) of a system.The sum of all kinetic and

potential energy of all components of a system

For molecules in a chemical system, the internal energy would include: the motion and interactions of the

molecules the motion and interactions of the

nuclei and electrons found in the molecules


Internal Energy:Extensive property

depends on mass of system

Influenced by temperature and pressure

Has a fixed value for a given set of conditions

State function


The internal energy of a system is a state function.A property of the system that is

determined by specifying its condition or its state in terms of T, P, location, etc

Depends only on its present condition

Does not depend on how the system got to that state/condition


The internal energy of a system can change when:heat is gained from or lost to the

surroundings work is done on or by the system.

The change in the internal energyE = Efinal - Einitial

E = change in internal energyEfinal = final energy of systemEinitial = initial energy of system


If Efinal > Einitial,

E >0 (positive) the system has gained energy

from the surroundings.

endergonic


The decomposition of water is endergonic (E > 0):

2 H2O (l) 2 H2 (g) + O2 (g)

H2 (g), O2 (g)

H2O (l)

E

Energy must be gained from the

surroundings.

final

initial


If Efinal < Einitial,

E < 0 (negative) the system has lost energy to

the surroundings.

exergonic


The synthesis of water is exergonic (E < 0)

2 H2 (g) + O2 (g) 2 H2O (l)

H2 (g), O2 (g)

H2O (l)

E

Energy is lost to the

surroundings in this

reaction.

initial

final


The internal energy of a system can change when energy is exchanged between the system and the surroundingsHeatWork

The change in internal energy that occurs can be found:

E = q + w

Where q = heatw = work


By convention:q = positive

Heat added to the systemw = positive

Work done on the system by the surroundings

q = negativeHeat lost by the system

w = negativeWork done by the system on the surroundings


Example: 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.

Given: system absorbs 140. J heat =system does 85 J work =

Find: E

+ 140. J- 85J


E = q + w

E = +140 J + (-85 J)

E = +55 J

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Unit 6: Thermochemistry