Unit 1: Thermochemistry

Unit 1: Thermochemistry

Introduction The First Law of

Thermodynamics Enthalpy Enthalpy of Reaction Calorimetry Hess’s Law Enthalpy of Formation

Introduction Most daily activities involve

processes that either use or produce energy:

Activities that produce energyMetabolism of foodBurning fossil fuels

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

Introduction Thermodynamics

The 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

Attractive and repulsive forces include:Gravity

Electrostatic forces between charged particlese- has potential energy due to its position near the positively charged nucleus

Most important attractive/replusive forces in chemistry

Introduction Attractive and repulsive forces

within a substance lead to a type of potential energy called chemical energy

The potential energy stored in substances resulting from the arrangements of the atoms in the substance

Introduction Units of Energy

SI 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 On the exam, you must be able to

convert from one set of energy units to the other using dimensional analysis.

You must know the conversion factors given on the previous slides!!!

Introduction

Example: Convert 725 cal to kJ.

Introduction

Example: A particular furnace produces 9.0 x 104 BTU/hr of heat. If 1.00 BTU = 251.9958 cal, use dimensional analysis to calculate the number of kJ of heat delivered by the furnace after running for 2.50 hours.

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)EEnergy 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 system w = 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 240. J of heat from the surroundings and does 85 J of work on the surroundings.


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Unit 1: Thermochemistry