Chem Volume 2

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Chapter 5: Heat Effects and Energy Relationships in Chemical Reactions

EPISODE 16: ENERGY CONSERVATION

OVERVIEW

The previous lessons were about the different forms of matter and the changes of

matter undergoes. This episode is the first of four lessons on energy relationships

accompanying changes in matter. We will begin with ideas on energy conservation

and energy transformation when matter undergoes change.

OBJECTIVES

At the end of this lesson, the student should be able to:

1. define basic concepts in thermodynamics energy, thermodynamics, system,surroundings;

2. differentiate open, closed and isolated system;3. explain the concepts of heat and work as energy transferred;4. state the First Law of Thermodynamics;5. give the operational definition of internal energy;6. explain the consequence of heat absorption and performance of work oninternal

energy;

7. explain the internal energy as a state function; and8. explain the energy transformations that accompany common life processes.

INTEGRATION WITH OTHER LEARNING AREAS

When matter undergoes change, it is always accompanied by a change in energy

content. In Episode 10, one of the indicators of chemical change is absorption or

release of heat. Physical changes described in Episode 7 Phases of Matter are also

accompanied by energy changes. Quantitative aspects of energy transformations areimportant topics in physics and engineering. In our homes, energy is an important

issue from the food we eat, the activities we do, the materials that we use all

produce or use up energy.

SCIENCE AND HEALTH IDEAS

Energy in food and nutrition Mans use of energy

SCIENCE PROCESSES

Observing

Experimenting

Classifying

Inferring


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Using mathematical relationships Predicting

VALUES

Responsibility for one's actions

LIFE SKILL

Decision making

IMPORTANT CONCEPTS

1. Thermodynamics is the study of energy changes accompanying physical andchemical changes that matter undergoes.

2. The First Law of Thermodynamics states that energy in the universe isconserved.

3. Energyis the capacity to do work.4. Heat and work are forms of energy in transit.5. Heat capacityis the amount of heat that has to be absorbed by the system for its

temperature to rise by one degree Celsius.

6. The total energy present in a given sample of matter is called its internalenergy.

7. Internal energy is a state function.8. Energy can be transformedfrom one form to another.BACKGROUND INFORMATION/EPISODE CONTENT

Energy is a word that almost everyone reads or uses frequently, but many of us just

have a vague understanding of its meaning. But we know it is important to life, we

use it, we produce it, and it can take many forms. What is energy and what forms

can it take?

Thermodynamics. The study of energy and its transformations is called

thermodynamics. This area of study straddles both physics and chemistry. Energy

conversions to mechanical forms are important concerns in physics. On the otherhand, chemical thermodynamics deals with the energy changes accompanying

physical and chemical changes that matter undergoes. The succeeding sections


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introduce us to some important terms and concepts used in describing

thermodynamic properties.

System and Surroundings. In order to study how energy changes during physicalor chemical change, we have to first divide the universe into two parts: the system

and the surroundings. A system is a portion of the universe that we isolate for the

purpose of study. It is a definite quantity of matter undergoing a change. The system

can be the reactants and products of a reaction, or a piece of ice melting on the

tabletop, or our bodies. It is separated from the rest of the universe by a real or

imaginary boundary. The beaker or flask where the reaction takes place denotes a

real boundary. However, the parts of the melting ice that do not touch the table have

an imaginary boundary.

Everything outside the system is the surroundings including boundaries. Thus, the

beaker where a reaction is happening is part of the surroundings of the reactants and

products of the reaction. The relationship between the components of the universe is

given by the equation:

Universe = System + Surroundings

Types of Systems. Systems may be classified according to the nature of its

interaction with its surroundings into three types:

In an open system, both energy and matter can be exchanged between thesystem and its surroundings. An example of this is a hot cup of coffee as it cools.It loses heat to the surroundings and loses water vapor (matter) in the form of

steam.

In a closed system, energy can be exchanged between the system and itssurroundings but matter cannot be exchanged. A stoppered glass flask of coffee

loses heat to its surroundings but does not lose water vapor to it.

In an isolated system, neither energy nor matter can be exchanged between thesystem and its surroundings. Hot water placed in a thermos bottle approximates

an isolated system. It does not readily lose energy to and neither does it losewater to its surroundings. However, the thermos bottle is not a perfect boundary

for an isolated system since energy is slowly lost to the surroundings, so that in

time, the water inside cools.

Energy, Work, and Heat. One of the most important knowledge that we have

about energy is that it is conserved. In other words, the energy of the universe is

constant. This idea is essentially the First Law of Thermodynamics, or more

familiarly, the Law of Conservation of Energy, which states that energy can

neither be created nor destroyed, but can be converted from one form to another.

Energyis often defined as the capacity to do work.Work, on the other hand, is done

when force is exerted over a distance. Whenever we define work, a situation


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showing work is being done is always cited. A crawling baby, a boy climbing a tree,

and a carabaoplowing the fields all show that work is being done. Moving objects

do work when they slow down or increase in speed. Maybe not as easily evident as

in the other examples, work is also being done when a gas like helium expands or is

made to occupy a larger space, or when it is compressed to a smaller space. Theseevents all show that energy is being used for them to occur.

Heatis the energy transferred between a system and its surroundings as a result of a

temperature difference. Energy, as heat, passes from a warmer body (with higher

temperature) to a colder body (with lower temperature).Like work, heat is energy in

transitbetween the system and its surroundings. It is a form by which energy can be

transferred across a boundary between the system and its surroundings.

Heat and work are the forms by which energy gets transferred or converted to other

forms. They can also be transformed into one another. Heat can be used to do work,

such as in steam engines. On the other hand, if you rub your hands together and your

hands begin to feel warm, work is being converted to heat.

Units of Energy.Both heat and work can be expressed in the same unit of measure,

typically thejoule, J. One joule, 1 J, is approximately the amount of energy needed

to lift a 1 kg object about 4 inches against gravity. The energy required by the

average human heart to make one beat is also about one joule. Another unit that is

often used to describe heat and work is the calorie, cal. One calorie,1 cal, is the

amount of heat needed to raise the temperature of one gram of water by 1oC. Onecalorie is equal to 4.184 J.

In food labels, the energy equivalent of food is expressed in big calories. When

calorie is abbreviated with a capital C, it means kilocalories or 1000 calories. The

label of a certain butter states that 15 g of that butter has an energy equivalent of 110

Calories or 110,000 calories or 460,240 joules!

The amount of work or heat involved in a physical or chemical change can be

calculated using some mathematical relationships arising from the definitions ofthese forms of energy and the first law of thermodynamics. Some simple examples

are given in the following sections.

Work.Work is often defined as the product of force and the distance through which

the force acts as described in the equation below. The equation is useful in solving

problems in physics that ask how much work is done when a car moves, or a ball

rolls downhill or when we climb the stairs.

Work (w) = force (F) x distance (d)


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But let us use an example which has some chemistry in it, like the expansion or

compression of a gas. During gas expansion or compression, the force, F, is equal to

the pressure and the area on which this pressure is exerted. The equation becomes:

Work (w) = (pressure x area) x distance (d)

But (area x distance) is actually the change in volume of the gas or the container

during the expansion or compression. Thus, work is measured by the pressure

multiplied by the change in volume of the gas.

Work (w) = V x P

When a gas expands, force is exerted on the surroundings, since the system increases

in volume and moves against the external pressure.

V is positive (final volume isgreater than initial volume) but work here carries a negative sign since work is

directed out of the system. On the other hand, when a gas is compressed, force is

exerted on the gas, causing its volume to decrease. V is negative, but work ispositive. Hence, we refine the equation to:

Work (w) = - (V x P)

Example: Calculate the work done when a gas expands from 40 liters to 58 liters at

a constant external pressure of 10 atm or a gas at constant pressure.

w= - (V x P)

w= - (18 L x 10 atm) = - 180 L-atm

To convert to the energy unit, joule, J, the conversion factor is 1 L-atm = 101 J.

w= - 180 L-atm x 101 J/L-atm = -18180 J or -18.2 kJ

Heat.When two objects are in contact, heat always flows from the object of higher

temperature to the one of lower temperature.If a cold metal spoon is placed in a cup

of hot coffee, heat will flow from the hot coffee to the metal spoon until their

temperatures are the same. How much heat is absorbed by the spoon? How much

heat is lost by the hot coffee?

Heat is represented by the symbol q. In interactions between a system and its

surroundings, the amount of heat gained by the system is the same amount of heatlost by the surroundings and vice-versa. Hence, the sum of the heat absorbed and


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heat evolved is equal to zero.

qsystem+ qsurr= 0

An increase in the temperature of a system implies that the system has absorbed or

gained heat. This is indicated by a positive value of q. When the temperature of the

system decreases, it evolves or loses heat to the surroundings and is indicated by anegative value of q.However, the magnitude of temperature change is dependent on

the amount of heat absorbed or lost and the heat capacity of the system, a property

that is characteristic of the substance or material that makes up the system.

Heat Capacity. Heat capacity is the quantity of heat needed to change the

temperature of a given object by one degree Celsius. If the system is made of the

same material throughout, or of a single substance, a more useful way of expressing

heat capacity is the specific heat capacity of the substance. Specific heat capacityis

the quantity of heat required to raise the temperature of one gram of the substance

by one degree Celsius. Water has a specific heat of 4.184 J/g-oC. The mathematical

equation used to determine how much heat has been absorbed or lost by a system is:

Q = heat capacity x T

or

Q = specific heat x mass of system x T

Example: How much heat must be absorbed by a 42 g piece of iron, Fe, to raise its

temperature from 22oC to 80oC? The specific heat of iron is 0.45 J/g-oC.

Q = (0.45 J/g-oC) (42 g) (80

oC 22

oC)

Q = 1096 J

The value of heat obtained in the example has a positive sign, indicating that heat

flowed into the system which is the piece of iron metal. Can you tell how much heat

was lost by the surroundings?

In the earlier section, we found that the heat gained by the system is exactly thesame quantity lost by the surroundings, such that the sum of these two quantities is

zero. Hence, the heat lost by the surroundings is also 1096 J, but this quantity carries

a negative sign.

qsystem + qsurr = 0

1096 J + (-1096 J) = 0

Internal Energy. Heat and work are the means by which the system exchanges

energy with its surroundings and they exist only during a change. A system does not

contain energy in the form of heat or work. The energy contained within the system,called internal energy, U, is the total energy, potential and kinetic, in various forms


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A property that is determined by the current state of the system and is not

determined by the manner in which the system achieved it is called a state function.

Internal energy, therefore, is a state function.

This can be illustrated by a person going up a fourstory building. The initial state

(state 1) is at the ground level and the final state, state 2, is at the fourth level. Theperson has an initial internal energy at state 1 and a final internal energy at state 2.

How can the person get to the fourth floor and achieve the final internal energy? He

can go up the building by taking the elevator or taking the stairs. He can go straight

up or take the elevator to the third level, go down to the second level, and climb the

stairs to the fourth level. He would have achieved the final internal energy state

when he is at the fourth level, in whatever way he reaches it.

Energy Transformations.Our main source of energy for our daily activities is the

combustion of fuels coal, fossil fuels, alcohol, biomass, to name some. But our

need for energy is no longer just for keeping warm or cooking food. Today, most of

our energy needs are for machines to do work. The heat generated from the burning

of fuels must be converted to other forms of energy.

What forms can energy take? Energy can be classified into two general forms:

potential and kinetic. The energy possessed by a moving body is called kinetic

energy. Potential energyis energy due to a condition, position or composition, and

is associated with forces of attraction or repulsion between objects. There are more

specific forms that energy can be converted into: chemical, electrical, mechanical,

light, and heat.

The human body converts the chemical energy from the food we eat to heat, thus

keeping our body at the right temperature and into mechanical energy that enables

our bodies to move. The video lesson also traces the transformation of chemical

energy from fossil fuels into mechanical energy and into electrical energy that is

transmitted to homes and industrial plants where it is further converted into various

energy forms such as light, sound and heat. Can you describe the transformation of

energy when a car moves?

If the First Law of Thermodynamics is true, then why do we always hear that our

energy resources are running out? What we are really saying here is that there are

forms of energy that are more useful than others. Our main sources of energy are the

fuels that we burn, converting their internal energy into heat. But aside from cooking

our food and keeping homes in temperate countries warm, heat is not a very useful

form of energy. It has to be converted to work or to other forms such as mechanical

energy to run machines and engines or to electrical energy. Man has not been

successful in finding an efficient way to convert most of the heat from burning fuels

to more useful forms, so most of the heat is simply used to increase the temperature

of the surroundings, which is for us just a waste of energy.


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Guide Questions/Answers

1. What is thermodynamics?Thermodynamics is the study of changes in energy that accompany changes

that matter undergoes.

2. Differentiate the system from the surroundings.The system is a small part of the universe an object, a chemical reaction, a

sample of gas that is being studied. The remaining part of the universe

outside of the system is the surroundings.

3. Describe the three kinds of system mentioned in the video lesson.The three kinds of systems presented are the following:

open system one that can exchange matter and energy with thesurroundings;

closed system one that can exchange energy but not matter, with thesurroundings.

isolated system one that cannot exchange matter nor energy with thesurroundings

4. Differentiate potential from kinetic energy.Potential energy is energy due to the position of the sample relative to other

particles that interact with it, while kinetic energy is energy due to the

motion of the object.

5. What is heat? What is work?Heat is energy that transfers between two or more bodies due to differences

in temperature, while work is the energy transformed when a force applied

on a body causes movement or changes in position.

6. When a gas expands, is work being done on the system or on thesurroundings? Explain your answer.

Work is done on the surroundings when a gas expands. During expansion,

the volume of the gas increases and exerts force outward as its boundariesmove intothe surroundings.

7. State the First Law of Thermodynamics.The First Law of Thermodynamics states that energy cannot be created nor

destroyed, and it is conserved, but it can be converted from one form to

another.

8. Define heat capacity.Heat capacity is the amount of heat needed to raise the temperature of the

system or an object by one degree Celsius.


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VIEWING ACTIVITIES

Activities that show the use of energy can be viewed at segments 2:30 2:50.

The First Law of Thermodynamics is illustrated in segments 14:05 to 15:22.

POSTVIEWING ACTIVITY

Discuss the answers to the Guide Questions.

TEACHING TIP

Discuss with the students how the First Law of Thermodynamics is applied in the

various situations viewed in the lesson.

ASSESSMENT

Quiz. Choose the letter corresponding to the best answer.

1. A solution of BaCl2is added to a beaker that contains a solution of Na2SO4, themixture is heated to promote the formation of the precipitate, BaSO4. What kind

of system is illustrated in this example?

A. Open C. Isolated

B. Closed D. Non-ideal

2. Which of the following is correct about no. 1?A. System = BaCl2solution + beaker

B. System = BaCl2+ Na2SO4solutions + beaker

C. System = BaCl2+ Na2SO4solutions + beaker + burner used for heating

D. System = everything involved in the experiment + experimenter

3. In what kind of system is the energy of the system constant?A. Open C. Isolated

B. Closed D. Ideal

4. Which of the following is NOT a state property?A. Height C. Temperature

B. Velocity D. Mass

5. A gas expands as it absorbs 500 J of heat and performs 200 J of work. Whathappens to the internal energy of the gas?

A. It increases by 300 J. C. It increases by 700 J.

B. It decreases by 300 J. D. It decreases by 700 J.


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REFERENCES

American Chemical Society. (1997). Chemistry in context. (2nd

ed.). USA: McGraw-

Hill, Inc.

Brown, T. L., LeMay, H. E. & B. E. Bursten. (2004). Chemistry: the central

science.NJ: Prentice Hall International, Inc.

Silberberg, M. S. (2000). Chemistry: the molecular nature of matter and change.

NY: McGraw Hill, Inc.

Useful websites

http://en.wikipedia.org/wiki/Thermodynamics

http://www.shodor.org/unchem/advanced/thermo/
http://en.wikipedia.org/wiki/Thermodynamicshttp://www.shodor.org/unchem/advanced/thermo/http://www.shodor.org/unchem/advanced/thermo/http://en.wikipedia.org/wiki/Thermodynamics


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EPISODE 17: ENERGY AND CHANGE

OVERVIEW

Physical and chemical changes are accompanied by changes in energy content of thesystem. In many instances, the energy changes are measurable. One way of

measuring the change in the heat content that accompanies a process or change isthrough calorimetry, which is described in this lesson.

OBJECTIVES

At the end of this lesson, the student should be able to:1. define enthalpy, and change in enthalpy;2. describe what standard state means for the different physical states of matter;3. differentiate exothermic from endothermic change;4. define calorimetry;5. describe the determination of enthalpy change through calorimetry;6. construct and calibrate a coffee cup calorimeter; and7. determine the heat of reaction in an acid-base neutralization.


The lessons dealing with thermodynamics are covered in Episodes 16 Energy

Conservation, 17 Energy and Change, 18 Thermochemical Equations and the

Direction of Change, and 19 The Laws of Disorder. The changes in heat content

accompanying processes are expressed in units involving molar and stoichiometric

quantities of substances. The lessons in Episodes 11 The Mole and 12 Patterns of

Change are particularly useful. The energy changes in reactions also influence thedirection of adjustments of equilibrium systems when disturbed as discussed in

Episode 22 Changes in Equilibrium System: Le Chateliers Principle.

SCIENCE AND HEALTH IDEAS

Energy changes during chemical reactions

SCIENCE PROCESSES

Experimenting

ObservingUsing mathematical relationships

Classifying

InferringPredicting


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VALUES

Responsibility for one's actions

LIFE SKILL

Decision making

IMPORTANT CONCEPTS

1. Most chemical reactions absorb or release energy usually in the form of heat.Reactions that release energy are exothermicreactions, while those that requireenergy are called endothermicreactions.

2. The heat content of a substance is called its enthalpy, H.3. The change in enthalpy, H, is the change in heat content accompanying a

reaction or a physical change occurring at constant pressure.

4. The change in enthalpy of a reaction is the difference between the sum of theenthalpies of products and the sum of the enthalpies of the reactants.

5. The standard enthalpy of formationof a compound is the heat of reaction forthe formation of one mole of the compound from its constituent elements at their

standard states.

6. Calorimetry is the science of measuring heat changes accompanying physicaland chemical changes.

BACKGROUND INFORMATION/EPISODE CONTENT

One of the indicators of chemical changes described in Episode 10 Indicators ofChemical Change is the release or absorption of heat. In Episode 9 Condensed

Phases of Matter, we learned that energy change is also involved in physical

processes such as evaporation. The heat absorbed or evolved during physical or

chemical change is usually measurable.

Exothermic and Endothermic Reactions. One of the most familiar reactions that

involve a release of energy is burning or combustion of fuels. The usefulness of

combustion reactions is the large amount of energy they produce. Reactions thatproduce energy are called exothermic reactions. In exothermic reactions, the total

heat content of the products is less than the total heat content of the reactants and thedifference is the amount of energy released. On the other hand, reactions that require


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energy in order to happen are endothermic reactions. In endothermic reactions, the

total heat content of products is greater than the total heat of reactants. The reactions

involved in cooking, such as in the hard boiling of an egg, are endothermic

reactions.

Processes involving phase changes can be exothermic or endothermic. Evaporation

is an endothermic process. Condensation of gases into liquids, the opposite processof evaporation, is exothermic. The formation of solutions can be exothermic or

endothermic. When sodium hydroxide is dissolved in water, heat is released and anincrease in temperature of the solution is observed. On the other hand, the formation

of an ammonium chloride solution is an endothermic process.

Internal Energy. In Episode 16 Energy Conservation, internal energy, U, was

defined the energy associated with the motion of particles (kinetic energy) and the

intermolecular forces of attraction that hold the particles in solids and liquids

(potential energy). According to the First Law of Thermodynamics, a change in

internal energy, DU, involves heat and work, the two forms by which energy is

transferred, and is given by the equation:

DU = q + w

where DU = change in internal energy

q = heat absorbed or lost by the systemw = work done by the system on the surroundings during the change,

which

is also equal to (PDV).

Enthalpy.Enthalpy, H, is the heat content of a system. It is defined in

thermodynamic terms as:

H = U + PV

where U is the internal energy of the system and PV is the product of pressure and

the volume of the system. The change in enthalpy, DH is:

DH = DU + PDV

Substitution of the term for internal energy, we get:

DH = DU + PDV = (q+ w) + PDV = (q PDV) + PDV = q


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This implies that when pressure remains constant and with work limited to pressure

volume work only, the change in enthalpy is equal to the change in heat content.

DH = q

where q is the heat absorbed or lost by the systemduring the process.

Like internal energy, absolute values of heat content of substances are impossible to

measure, but changes in heat content during a chemical reaction are measurable.And again like internal energy, enthalpy is a state function, meaning that the change

in enthalpy, DH, is the difference between the enthalpies of the final state and the

initial state:

DH = Hf - Hi

where Hf and Hi are enthalpies of the final and initial states, respectively. This

implies that:

for an exothermic process, where heat is released during the reaction, theproducts have less heat content than the reactants. Therefore, Hf < Hi,

and

H < 0 or a negative value. For an endothermic process, where heat is absorbed during the reaction, the

products have higher heat content than the reactants. Therefore, Hf > Hi,

and

H > 0 or a positive value.

The change in enthalpy of a reaction is also referred to as heat of reaction, or

enthalpy of reaction or simply enthalpy change. If the reaction is a combustionprocess, the enthalpy change is also called heat of combustion or enthalpy of

combustion. If the process involves formation of a solution, the terms used maybeheat of solution or enthalpy of solution.

Determination of Enthalpy Change. Enthalpy changes accompanying processes

can be determined in several ways, and the choice of method depends on its

suitability based on the nature of the process. The direct way of measuring heat of

reaction is through calorimetry. Indirect ways involve the use of knownthermochemical data, such as the enthalpy of formation of substances. The

determination of the heat of reaction for the formation of carbon monoxide isdescribed below.

The heat of reaction for the formation of carbon monoxide, CO, is 110.88 kJ.


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C(s)+ O2(g)CO(g)H = 110.88 kJ

The reaction, as indicated by the negative value for H is exothermic and 110.88 kJ

of heat is evolved when 1 mole of solid carbon burns in oxygen gas to form 1 moleof carbon monoxide gas.

The enthalpy of reaction is defined as the difference between the enthalpies of the

final and initial states of the system,

DH = Hf - Hi

For chemical reactions, the initial state is when reactants have not yet reacted, andthe final state is when the products have been formed. Thus, Hfis also the sum of the

enthalpies of the products, and Hi is the sum of the enthalpies of the reactants, and

DH is the difference between the sum of enthalpies of the products (SHP) and the

sum of the enthalpies of the reactants (SHR):

H = SHP SHR

The enthalpy for each substance is the standard enthalpy of formation (Hf). The

standard enthalpy of formation of any substance is the heat of reaction for theformation of one mole of the substance from its constituent elements at their

standard states. This definition implies that only compounds can have nonzerovalues for standard enthalpies of formation.

What does standard state of an element mean? The standard state of an elementis

its most stable form at 1 atm and 25C. Solids and liquids are in their standard states

when they are in their pure forms. Gases are in their standard states when their

pressure is 1 atm. Solutions are in their standard states when their concentration is 1

M (one molar).

As an example, let us calculate the enthalpy of combustion of propane, given the

standard enthalpies of formation of the substances involved.

Standard enthalpies of formation, Hf:C3H8(g) -103.8 kJ/mol

CO2(g) -393.5 kJ/mol

H2O(g) -241.8 kJ/mol


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q = mSt = mS (tfinal tinitial)

= (100 g x 4.184 J/g-C)(50C 25C)

= 10,460 J x 1 kJ/(1000 J)q = 10.46 kJ

Calibrating the Coffee Cup Calorimeter. In the video lesson, a simplified

calorimeter, called the coffee cup calorimeter is constructed using two Styrofoam

cups with a cover, a stirrer and a laboratory thermometer. The device is calibrated to

determine its heat capacity. The calibration procedure involved the following steps:

1. Equal volumes of cold and hot water, whose temperatures were recorded, aremeasured.

2. These are mixed in the calorimeter.3. The equilibrium temperature (the highest temperature) of the mixture is

measured.

Assuming that the calorimeter is an isolated system, and no heat is exchanged with

the surroundings, then the final temperature recorded reflects the heat transfer that

occurred from the hot water to the cold water and the calorimeter. The calculation of

the heat capacity of the calorimeter was carried out as follows:

qsystem= qcalorimeter + qcold water + qhot water = 0

where qcalorimeter, qcold water and qhot water refer to the heat absorbed or released by thecalorimeter, cold water and hot water, respectively. Hence,

qcalorimeter + qcold water = qhot water

HC(tf tc) + mcS(tf tc) = mhS(tf th)

where HC = heat capacity of the calorimetertf = final temperature of the system

tc = initial temperature of the cold water

th = initial temperature of the hot watermc = mass of cold water

mh = mass of hot water

Measuring Heat of Reaction. The calibrated calorimeter was used to measure the

heat of reaction of an acidbase reaction. The reaction was the neutralizationreaction between hydrochloric acid (HCl) and sodium hydroxide (NaOH).

HCl(aq) + NaOH(aq) H2O(l) + NaCl(aq)

In the demonstration, equal volumes of 2 M HCl solution and 2 M NaOH solution

were used.


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1. The initial temperatures of the two solutions were measured separately.2. Twenty five, 25, mL of the NaOH solution was placed in the coffee cup

calorimeter, then 25 mL of the HCl solution was added.

3. The calorimeter was covered, the solution was stirred and the equilibrium orfinal temperature of the mixture was measured.

Since dilute solutions of acid and base were used, the specific heat (4.184 J/g-C)

and density (1.0 g/mL) were assumed to be the same as that of pure water. It was

also assumed that no heat was lost to the surroundings or absorbed from the

surroundings.

The enthalpy of the reaction was determined as follows:

qcalorimeter+ qsolution+ qrxn= 0

HC(tf ti) + mS(tf ti) + DHrxn= 0

where HC = heat capacity of the calorimeter

tf = final temperature of the calorimeter and its contents

ti = initial temperature of the solutions before mixing

m = total mass of the acid and base solutions

S = specific heat of the solutionsDHrxn = enthalpy of the reaction

The total mass of the solution is obtained by multiplying the total volume, 50 mL,

by the density of the solution, which is assumed to be the same as that of water.

m= dx V= (1.0 g/mL x 50 mL) = 50 g

The DHrxn obtained is for the actual amounts of reactants used in the acidbase

reaction. How much was used in the experiment? Using our knowledge onconcentrations of solutions described in Episode 13, the number of moles of HCland NaOH used are obtained by multiplying the concentration of the solution (2 M)

by the volume of the solution (25 mLconvert to 0.025 L) of the base or the acid

since the concentration and the volumes used were identical.

Number of moles of HCl (or NaOH) = Mx V= 2 mol /L x 0.025 L = 0.050 mole

The enthalpy change calculated is for 0.050 mol of HCl (or NaOH). You mayproceed to calculate the molar heat of neutralization of HCl and NaOH by dividing

the DHrxn obtained by 0.050 mol.


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VOCABULARY WORDS

1. Enthalpy heatcontent of a substance.2. Enthalpy of reaction the change in heat content of a chemical reaction

occurring at constant pressure.

3. Enthalpy of formation the heat of reaction for the formation of a compoundfrom its constituent elements at their standard states.

4. Specific heat the amount of heat needed to raise the temperature of one gramof substance by one degree Celsius.

5. Heat capacity the amount of heat needed to raise the temperature of a givenquantity of substance or an object by one degree Celsius.

6. Calorimetry the science of measurement of heat changes.7. Calorimeter device used in measuring heat changes accompanying chemical

reactions and physical changes.

PREVIEWING ACTIVITIES

A. Ask students to recall reactions that they have observed and classify thereactions as endothermic or exothermic.

B. Pose the Guide Questions which the students will answer after viewing the

episode. Ask them to focus on finding the answers to the guide questions as theywatch the video.


1.

In the construction of the coffee cup calorimeter, why are two styrofoamcups used instead of just one?

The second styrofoam cup provides another layer of insulation from the

surroundings, creating an isolated system that does not exchange matter and

heat to the surroundings. The change in temperature recorded in the

calorimeter would be more accurate.

2. What characteristic(s) of styrofoam makes it a good material for thecalorimeter?

Styrofoam is a light plastic material used to make insulation and packaging

materials. It contains very small packets of air that does not allow rapid heat

transfer since air is not a good thermal conductor. Heat from the inside of

the calorimeter will take longer time to escape to the surroundings.


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3. What is the difference between specific heat and heat capacity?The specific heat, S, is the amount of heat needed to raise the temperature of

one gram of a substance by one degree Celsius. On the other hand, the heat

capacity, C, is the amount of heat needed to raise the temperature of a given

quantity of substance by one degree Celsius.

4. What does it mean when a substance has a high specific heat?A substance with high specific heat requires a large amount of energy to

raise its temperature. The substance will not easily become hot. Water has a

relatively high specific heat at 4.184 J/g-oC. In comparison, metals have low

specific heat values. Mercury, for instance, has a specific heat of only 0.138

J/g-oC. It will require 30 times more heat to raise the temperature by 1

oC of

a 1-g sample of water compared to 1 g of mercury. If 5 J of heat is applied

to each of 1-g samples of water and mercury, the temperature of the water

sample will be higher by 1.19oC but the mercury sample will be hotter by

36.23oC!

5. Why do we need to calibrate the coffee cup calorimeter?During heat transfers, the calorimeter itself will absorb or lose some amount

of heat. The calibration is essentially the determination of the heat capacity

of the calorimeter, which should be taken into consideration for a more

accurate determination of the enthalpy of reactions or processes made to

occur in the calorimeter.

6. What are the possible sources of error during the calibration?Possible sources of error include:

a. not stirring the mixture wellb. not placing the cover tightlyc. pulling out the thermometer from the calorimeter when reading the

temperature

d. excluding the temperature of the calorimeter in the measurement of theinitial temperature

e. reading the temperature of liquids before measuring the amount needed,using a measuring glassware that has not been equilibrated into thetemperature of the liquid

VIEWING ACTIVITIES

In the absence of materials,

two segments cited here would be useful in the understanding of calorimetry:

(1)Calibration of the Coffee Cup Calorimeter at 14:28 19:13 and(2)Determination of the Heat of Reaction Between

an Acid and a Base at 20:50 - 22:50.


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POSTVIEWING ACTIVITIES

A. Discuss the answers to the Guide Questions.B.

Discuss the principle involved in calorimetry.

C. Let the students to construct and calibrate their own coffee cup calorimeter.D. Discuss the computations involved in Activity C.E. Let the students determine the enthalpy change for an assigned reaction.F. Discuss the computations involved in Activity E.ASSESSMENT

A. Quiz. Write the letter corresponding to the best answer.

1. How do you classify a reaction where the total energy of the products isgreater than the total energy of the reactants?

A. Endothermic C. SpontaneousB. Exothermic D. Nonspontaneous

2. What law is applied in calorimetry?A. Law of Conservation of Mass C. Law of Conservation of EnergyB. Law of Definite Composition D. Law of Multiple Proportions

3. Which is in its standard state?A. Br2(l) C. NaCl (1 M)B. O2(g, 1 atm) D. A, B and C

4. Which is CORRECT about the enthalpy change for a reaction?A. It is dependent on the amount of reactants.B. It is negative if the reactants release energy to the surroundings.C. It is the heat change measured at constant pressure.D. A, B and C.

5. Which of the following can absorb the greatest quantity of heat if equal

masses of are heated from 0C to 20C?

A. Aluminum (Al) metal, S = 0.900 J/g-CB. Graphite (C), S = 0.720 J/g-CC. Iron (Fe) metal, S = 0.444 J/g-CD. Gold (Au) metal, S = 0.129 J/g-C


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ANSWER KEY

A. Quiz.1. A 2. C 3. D 4. D 5. A

6. C 7. D 8. C 9. D 10. C

B. Problem Solving.1. Since it is assumed that there is no heat lost,

heat gained by water + heat lost by aluminum = 0

heat gained by water = heat lost be aluminum

mwaterx Swaterx (tf ti)water= mAlx SAlx (tf ti)Al

100 g x 4.184 J/g-C x (tf 20

C) = 50 g x 0.900 J/g-

C x (tf 60

)

tf= 23.9C

2. There is no heat lost, therefore:

heat released by coal = heat absorbed by water

mcoalx DHcoal= mwaterx Swaterx (tf,water ti,water)

mcoalx 30,000 J/g = 1,000 g x 4.184 J/g-C x (100C 25C)

mcoal= 10.5 g

REFERENCES

Brown, T. L., LeMay, H. E. & B. E. Bursten. (2004). Chemistry: the central science.

NJ: Prentice Hall, Inc.

Chang, R. (2005). Chemistry. NY: McGraw-Hill, Inc.

Silberberg, M. S. (2000). Chemistry: the molecular nature of matter and change.

NY: McGraw-Hill, Inc.


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EPISODE 18: THERMOCHEM ICAL EQUATI ONS AND THE DIRECTI ONOF CHANGE

OVERVIEW

This episode is the second lesson on thermochemistry, the study of heat effects that

accompany chemical reactions. Writing thermochemical equations and calculating

the energy change that accompanies a process from known thermochemical data arethe main topics in this lesson. The relationship of enthalpy changes to other

thermodynamic properties that influence the spontaneity of a process is introduced.

OBJECTIVES


1. write thermochemical equations;2. define standard enthalpy of formation and standard enthalpy of reaction;3. solve for the standard heat of reaction (DHr) from standard heat of formation

(DHf) values;4. recognize the information about substances and chemical reactions from

enthalpy of reaction;

5. calculate the heat of reaction (DHr) using Hess Law;6. calculate the heat of combustion (DHc) of fuels; and7. compare spontaneous reactions and nonspontaneous reactions.


This episode is part of the big chapter on thermodynamics which began with

Episode 16 - Energy Conservation and runs up to Episode 19 - The Laws of

Disorder.

SCIENCE PROCESSES

ObservingExperimenting

Using mathematical relationships

InferringClassifying

Predicting

VALUES

Appreciation for the natural resources Care for the environment


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L IFE SKILLS

Judicious use of energy Conservation of energy sources

IMPORTANT CONCEPTS

1. Chemical and physical changes can be exothermic or endothermic.2. Information about the heat change accompanying a reaction is indicated in a

thermochemical equation.

3. The reverse of an exothermic process is endothermic.4. Enthalpy change can be determined indirectly from known thermochemical

data.

5. Enthalpyis a state function, and the value is the same regardless of the steps orroute taken to get from the initial to the final state. This is the basis of H essLaw, which states that the enthalpy change for a process is the same as the sumof the enthalpy changes of all steps that comprise the process.

6. Fuel efficiency can be compared from their heats of combustion, or from fuelvalues.


Thermochemistry is the quantitative study and measurement of heat changes thataccompany physical and chemical changes. The change in heat content during achemical reaction is called heat of reaction(qrxn), which is defined as the quantityof heat exchanged between a system and its surroundings when a chemical reaction

occurs within a system. The change in enthalpy, DH, is the difference between theenthalpies of the final state, Hf, and the initial state, Hi.

DH = Hf - Hi

When heat is released during a reaction, the products have less heat content thanthe reactants, thus H f < H i. Hence, H < 0 or a negative value, and thereaction is exothermic.

When heat is absorbed during the reaction, the products have higher heatcontent than the reactants, thus H f > H i. H > 0 or a positive value, and the

reaction is endothermic.


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The correct way of writing a thermochemical equation is shown below.

H2(g) + O2(g) H2O(l) DH25C= 285.8 kJ

where the superscript () indicates standard state and the subscript indicates the

temperature at which the value was measured. Ordinarily, however, DH values are

taken to be at 25o

C, unless otherwise specified.

Proper Use of Thermochemical Equations. Thermochemical equations can beused effectively by applying the following three basic rules of thermochemistry.

1. The magnitude of DH is directly proportional to the amount of reactant orproduct. This is something we know from experience, that heat content is anextensive property. The more gasoline we burn, the more energy we get!

This rule allows us to findD

H for any amount of reactant or product. Considerthe equation given as an example earlier:

H2(g) + Cl2(g) 2HCl(g) DH = 185 kJ

This equation means that 185 kJ of heat are produced when

one mole of H2reacts completely, or one mole of chlorine gas, Cl2, reacts completely, or two moles of hydrogen chloride are formed from H2and Cl2.Suppose only 1 mole of HCl was formed. What will be DH?

DH = 1.0 mol HCl xmolHCl

kJ

2

185- = -92.5 kJ

2. H for a reaction is equal in magni tude but opposite in sign to the H for thereverse reaction.This means that the amount of heat released from the formation

of HCl from H2and Cl2in the example is the same amount of energy needed tobe absorbed to decompose the same amount of HCl.

H2(g) + Cl2(g) 2HCl(g) DH = 185 kJ

2HCl(g) H2(g) + Cl2(g) DH = +185 kJ

A qualitative statement of this rule is that if a certain reaction is exothermic, thenits reverse is endothermic. In Episode 9 Condensed Phases of Matter, we

learned that the process of evaporation of a liquid is an endothermic process. The


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reverse process, condensation of a gas to a liquid, must be exothermic or energy

releasing.

3. The value of H for a reaction is the same whether the process occur s in asingle step or a series of steps.This principle is known as Hesss Law, which

states that if a reaction occurs in several stages or steps, the enthalpy change forthe overal l process is the sum of the enthalpy changes for the individual steps.

Let us consider the formation of nitrogen tetroxide, N2O4, from N2(g)and O2(g).

The reaction occurs through 2 steps:

Step 1: N2(g) + O2(g) NO2(g) DH= + 33.85 kJ

Step 2: 2 NO2(g) N2O4(g) DH= 58.04 kJ

The overall reaction is obtained by:

Step 1 x 2: N2(g)+ 2 O2(g) 2 NO2(g) DH= 2 (+33.85) kJ

Step 2: 2 NO2(g) N2O4(g) DH= 58.04 kJ

Since the overall reaction is obtained by adding the last 2 equations, the enthalpy

change for the overall reaction is obtained by adding the values for the 2

equations.

We get:

N2(g)+ 2 O2(g)N2O4(g) DH = (2 x 33.85) 58.04 kJ

DH = + 9.66 kJ

Notice that the use of Hess Law and knowledge of heats of reaction of somechemical reactions is also one way of determining the heat of a reaction.

What Information the Heat of Reaction Gives. The sign of DH reveals someinformation about the process itself. Reactions involve bond breaking, which is

energy requiring and bond formation, which is energy releasing. For example, the

last equation represent a reaction that is endothermic, in which heat is absorbed

during the formation of N2O4(g). The positive value of the overall DH shows that alarger amount of energy is needed to break bonds in the reactants than the amount of

energy released to form the bonds in the products. Since the amount of energyneeded is greater than the amount of energy released, the deficiency in energy is

taken from an external source, in some instances, from the surroundings.

The heat of reaction also indicates the relative stability of the products of a chemical


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reaction. The greater the amount of heat evolved during the formation of a

compound, the more stable is the compound. This means that exothermic reactions

form products that are more stable than the reactants.

Aside from predicting the stability of the products, it is also possible to predict thechemical reactivity of related compounds from their heats of reaction. For example,

the standard enthalpies of formation of the compounds HF, HCl, HBr, and HI are -271.1, -92.3, -36.4 and +26.5 kJ/mol respectively. Since the value for HF is most

exothermic, this indicates that HF is the least reactive and most stable among thecompounds, and would require the greatest amount of energy to decompose.

Determination of Enthalpies of Reaction. In Episode 16 Energy Conservation,the determination of the heat of reaction between HCl and NaOH through

calorimetry was demonstrated. Calorimetry is the experimental and direct way of

determining heats of reaction. However, many reactions cannot be carried out in

calorimeters and the enthalpy changes for these reactions are determined indirectly

from known thermochemical data.

One way of calculating heats of reactions is through the application of Hess Law, asshown in the section on the proper use of thermochemical equations. The enthalpy

change for reactions that are difficult to carry out, such as the formation of carbon

monoxide from the incomplete combustion of carbon, can be calculated using this

method. The heats of reaction for the complete combustion of carbon (Equation 1)

and the reaction of carbon monoxide with oxygen to produce carbon dioxide(Equation 2) are known and can be used to calculate the heat of formation of carbonmonoxide.

Equation 1: C(s) + O2(g) CO2(g) H = -393.5 kJ

Equation 2: CO(g) + O2(g) CO2(g) H = -283.0 kJ

To calculate H for the reaction: C(s) + O2(g) CO(g),

a. reverse equation 2: CO2(g) CO(g) + O2(g) H = + 283.0 kJb. add equation 1 to the reverse of equation 2.

CO2(g) CO(g) + O2(g) H = + 283.0 kJ

+ C(s) + O2(g) CO2(g) H = - 393.5 kJ

C(s)+O2(g)+ CO2(g) CO2(g)+ CO(g)+ O2(g) H = +283.0 + (-

393.5)kJ

Net equation: C(s) + O2(g) CO(g) H = -110.5 kJ


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Fuel Value. One of the important uses of thermochemical measurements is inassessing materials as energy sources. Most of these materials are what we call fuels

which liberate heat through combustion. One way to compare fuels is through theirheats of combustion. In general, fuels with higher heat of combustion are better

fuels. Among the fuels listed in the table below, we find octane, the maincomponent of gasoline, to have the highest heat of combustion.

Fuels can also be compared through their fuel values. Fuel value is the amount ofheat produced per gram instead of per mole of burned fuel. From the list below,

hydrogen gas yields the most amount of energy per unit mass.

Fuel ValuesFuel HC, kJ/mole Fuel value (kJ/g)

Hydrogen, H2 282 141

Methane, CH4 880 55

Methanol, CH3OH 726.4 22.7

Ethanol, C2H5OH 1366.7 29.7

Octane, C8H18 4947.6 43.4

Spontaneous and Nonspontaneous Reactions. Combustion reactions arereactions that, once started, can continue on as long as reactants are still available.

Chemical reactions and physical processes that occur without external intervention

or application of energy are called spontaneousprocesses. Examples of these arethe diffusion of a colored substance in water, burning a piece of paper, browning of

cut potatoes, and rusting of metals. On the other hand, processes that take place only

upon the application of heat, such as boiling an egg, are nonspontaneousreactions.

What makes a process spontaneous? If we look at some familiar processes that are

spontaneous, we notice that these processes tend to go in the direction of lower

energy state. A ball will roll downhill spontaneously accompanied by a drop inpotential and kinetic energy when it rolls to a stop. Combustion reactions are

exothermic reactions.

Many exothermic processes occur spontaneously but there are some which are not

spontaneous. The freezing of water is also an exothermic process, but water will notfreeze unless we place it in a freezer. But in Alaska, water will freeze

spontaneously! This tells us that temperature also influences the spontaneity ofcertain processes.


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There are also endothermic processes that occur spontaneously. One example of a

spontaneous process that is endothermic is the dissolution of ammonium chloride.

What really drives spontaneous processes to occur? That is the topic of the next

lesson.

VOCABULARY WORDS

1. Enthalpy the heat content of a substance or a system.

2. Standard heat of reaction change in heat content during a reaction wherereactants and products are in their standard states.

3. Standard enthalpy of formation a special case of standard heat of reaction,where the reaction involved is the formation of a compound from its component

elements in their standard states.

4. Thermochemistry the quantitative study and measurement of heat changesaccompanying chemical processes.

5. Thermochemical equation a chemical equation that includes the change inheat content that accompanies the processes being represented by the equation.

6. Heat of combustion enthalpy change accompanying the combustion of onemole of a substance.

7. Hesss Law law that states that the enthalpy change of the overall process isequal to the sum of the enthalpy changes of all the steps that comprise the

process.

8. Fuel value the heat liberated from the combustion of one (1) gram of fuel.9.

Spontaneous process a process that occurs without any intervention once thereactants come together at the right conditions.

10.Nonspontaneous reaction a process that requires a continuous input of energyin order to occur.

PRE-VIEWING ACTIVITIES

A. As a motivational activity, demonstrate to the students the reaction between

water and sodium metal. Place a very small piece of sodium metal in a dish with

a small amount of water. (WARNING!This should be done by a capable person


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in a proper venue, using the appropriate equipment and protective gear. The

audience should be in a safe distance away.) Ask the students to observe the

reaction carefully. The reaction of sodium metal with water is very explosive.

This is an example of an exothermic reaction wherein energy is liberated in the

form of heat and light.

B. Introduce the episode that deals with the concepts of enthalpy and enthalpychanges by asking how energy changes are expressed as part of chemical

equations.

C. Pose the Guide Questions that the students will answer after viewing the

episode. Ask them to focus on finding the answers to the guide questions as they

watch the video.


1. What are included in a thermochemical equation?A thermochemical equation includes a balanced equation according to theLaw of Conservation of Mass, the physical states of the reactants andproducts are indicated, the temperature at which the reaction occurs, and theheat of reaction carrying the appropriate sign.

2. What does the heat of reaction tell us?It indicates whether the reaction is endothermic or exothermic; whetherenergy for bond breaking is greater or less than the energy released frombond formation, and the relative stabi l i ty of the products for med.

3. How is fuel value calculated?

Fuel value is calculated by divi ding the molar heat of combustion of a fuel byi ts molar mass to get a value in kJ/g fuel .

4. How is the enthalpy change of a reaction that takes place in several steps

determined according to Hess Law?

According to Hess Law, the enthalpy change of a reaction that occurs inseveral steps can be calculated simply by adding the enthalpy changes for allthe steps that compr ise the reaction.

5. Why is hydrogen not used as a commercial fuel?

Hydrogen may have a very high fuel value but its being a gas with an

extremely low critical temperature makes it difficult to store and is lesspor table compared to other fuels.


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6. What is the relationship between enthalpy, entropy, and spontaneity of a

reaction?

The spontaneity of a reaction is influenced by both enthalpy and entropy.Many spontaneous reactions have negative enthalpies (exothermic) and

posi tive entropy changes (going to more disorder).

VIEWING ACTIVITIES

Let the students view the segments on(1)Writing Thermochemical Equations at 2:30 - 8:30 & 9:00 - 19:00 and

(2)Introduction to Entropy at 19:05 - 20:50.

POSTVIEWING ACTIVITI ES

A. Discuss the answers to the Guide Questions.

B. Discuss how the Fir st Law of Thermodynamics is applied as per the videoclip.

TEACHING TIP

Concept Mapping.

Divide the students into groups and ask them to prepare a concept map using thevocabulary words given previously by providing linking words to relate the

concepts. Ask the groups to present their concept maps in class.

ASSESSMENT

Quiz.

Give the required information.

1. What is thermochemistry?

2. Give two ways on how enthalpy of a reaction can be determined.

3. What refers to the degree of disorderliness?


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4. What law gives the relationship between heat, work and change in internal

energy?

5. What does the symbol PDV means?

6. The enthalpy of formation of the amino acid leucine, C6H13O2N(s) is -637.3

kJ/mol. Write the thermochemical equation to which this value applies.

7. Calculate the standard heat of formation for the fermentation of glucose:

C6H12O6(s) 2C2H5OH(l) + 2CO2(g)

8. Calculate DHo for the synthesis of lime, CaCO3, an important step in the

manufacture of cement, from the following thermochemical data:

CaCO3(s) CaO(s) + CO2(g)

DHof, CaCO3(s)= -1206.9 kJ/mol

DHof, CaO(s) = -635.1 kJ/mol

DHofCO2(g) = -393.5 kJ/mol

ANSWER KEY

Quiz.

1. Thermochemistry is the study and measurement of heat changes that take placeduring chemical reactions.

2. Two ways on how enthalpy of a reaction can be determined:a. through calori metry.b. fr om standar d enthal pies of for mation of substances.

3. Entropy

4. First Law of Thermodynamics

5. Pressur e-volume work, i .e. expansion or compression

6. C(s) + 13/2 H2(g) + O2(g) + N2(g) C6H13O2N(s) DHo

f= 637.3kJ/mol


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EPISODE 19: THE LAWS OF DISORDER

OVERVIEW

Things around us change even without human intervention. Iron nails rust, fruits

ripen, leaves fall from their branches. But other processes need some help to occur.

Food will not cook without heat. Water will decompose to hydrogen and oxygen

only if electricity is continuously applied through it.

This episode, the fourth in this volume on thermodynamics, is about the

thermodynamic property called entropy, which is related to the degree of

randomness of a system, and how the change in entropy influences the spontaneityof reactions.

OBJECTIVES


1. describe a spontaneous reaction;2. give example of spontaneous changes;3. define entropy in his/her own words;4. determine when a process results to an increase or a decrease in entropy;5. explain the Second Law of Thermodynamics;6. define free energy;7. explain the Third Law of Thermodynamics; and8. predict whether changes are spontaneous or not based on the change in free

energy (G).


This episode can be best appreciated after going through with the three previous

episodes in the same chapter, namely, Episode 16 Energy Conservation, 17 Energy and Change, and 18 Thermochemical Equations and Direction of Change.

SCIENCE PROCESSES

Observing

Experimenting


Classifying

Inferring

Predicting


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VALUES

Appreciation for order

LIFE SKILLS

Organization Orderliness

IMPORTANT CONCEPTS

1. A spontaneous process is a process that occurs by itself once started and willcontinue to happen without outside intervention.

2. A non-spontaneous processis an unnatural process that once started, requires acontinuous application of energy to proceed.

3. Entropy is a property of a system that is a measure of the randomness ordisorder of the system. The greater the degree of randomness or disorder in a

system, the greater is its entropy.

4. The Second Law of Thermodynamics states that every spontaneous changeproduces an increase in the entropy of the universe.

5.

The Third Law of Thermodynamics states that the entropy of a perfectcrystalline substance at absolute zero temperature is zero.

6. The free energy, G,is the energy available to do work.7. The enthalpy , H,of a system is the sum of the free energy, G, and the energy

used to increase the entropy of either the system or the surroundings, S.

8. The change in free energy accompanying a process, G, is given by the Gibbs

free energy equation,G = H TS.

9. If G for a process is negative, the process is spontaneous. If G is positive, the

process is nonspontaneous.


Spontaneous and Non-spontaneous Processes. Chemical and physical changes

can be either spontaneous or non-spontaneous. Spontaneous processes are those

that once started continue to happen without any need for outside intervention. A

ball rolling downhill is an example of a spontaneous mechanical change. Themelting of ice cubes in a glass of water at room temperature is a spontaneous


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physical change. A mixture of hydrogen gas and oxygen gas will explode in the

presence of a spark.

Nonspontaneous processesare those that do not take place without a continuous

outside intervention, driving force being applied on the system. A ball will rolluphill only if energy is used to roll the ball. Water will not decompose into hydrogen

and oxygen gases unless electricity is continuously applied.

When a process is spontaneous, its reverse is non-spontaneous in the same

conditions. A ball will roll spontaneously downhill, but an uphill roll is definitely

non-spontaneous. Dry ice or solid carbon dioxide will spontaneously sublime, but

solid carbon dioxide will not form unless vigorous conditions of extremely high

pressure and very low temperature are created.

Many spontaneous processes involve a decrease in energy. Combustion reactions are

exothermic and spontaneous. When a ball rolls downhill, its potential energy drops.

But what about the melting of ice? Or the dissolution of ammonium chloride,

NH4Cl, in water, both of which are endothermic? Besides enthalpy, there is another

driving force of spontaneous reactions, the randomness factor or entropy.

Entropy. Entropy, symbolized by S, is a property of a system that expresses the

degree of randomness or disorderliness of the system. The unit for entropy is J/mol-

K. In general, the more random or disorderly is the distribution of particles in a

system, the greater the entropy. All substances possess some degree of randomness

since particles are in constant motion and thus have positive entropy values.

The Third Law of Thermodynamics. Consider the two allotropes of carbon,

diamond and graphite, which have the same composition but have different physical

properties. Are their entropy values at the same conditions the same? Is entropy a

measurable quantity?

The determination of entropy values is made possible through the use of a reference

point where the entropy of a substance is zero. This condition is described in the

Third Law of Thermodynamics, which states that at absolute zero temperature, 0K, pure crystals of a substance are in a perfectly ordered state and thus have an

entropy value of zero. The lowest possible value for the entropy of a substance is

therefore zero, when the temperature is at absolute zero, 0 K.

When comparing entropies of substances or evaluating entropy changes

accompanying processes, the following conditions serve as useful guide:

1. The entropy of a substance always increases as it changes state from solid to

liquid to gas.

2. When a pure solid or liquid dissolves in a solvent, the entropy of the substance

increases.

3. When gas molecules escape from a solvent, entropy increases.


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and energy that is no longer usable to do work because it has been used to create

more disorder, S.

H = G + TS

On rearrangement, the equation becomes the Gibbs free energy equation,

G = H TS

where G = change in free energy, Gfinal - Ginitial

H = enthalpy change, Hfinal - HinitialS = entropy change, Sfinal - SinitialT = absolute temperature, K.

Spontaneous processes have a negative free energy change, G < 0. These

processes release energy to the surroundings and the greater and more negative G,

the greater is the likelihood of the process occurring. Non-spontaneous processeshave a positive free energy change, G > 0. For these processes to occur, the system

must take in energy from the surroundings. Without this additional energy, the non-

spontaneous process will not take place or will stop if it has started. If G = 0, then

the process and its reverse occur simultaneously, and the system is in equilibrium.

The following table summarizes the effects of enthalpy changes and entropy changes

on spontaneity, based on the Gibbs free energy equation:

Enthalpy change Entropy change Spontaneous? ExampleExothermic, H < 0 Increase, S > 0 Yes, G < 0 Oxidation of glucose

in muscle

Exothermic, H < 0 Decrease, S < 0 Only at low T,

if |TS|< |H|

Freezing of water

Endothermic, H > 0 Increase, S > 0 Only at high T,

if TS > H

Melting of ice

Endothermic, H > 0 Decrease, S < 0 No, G > 0

at all temperatures

Formation of C2H4

from C and H

VOCABULARY WORDS

1. Entropy the degree of randomness of a system.2. Second Law of Thermodynamics all processes in the universe tend to

proceed to a state of increasing entropy of the universe.

3. Spontaneous process one, which when started, will continue to happen evenwithout external intervention.

4. Non-spontaneous process one, which when started, will require continuous


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input of energy in order to continue happening.

5. Third Law of Thermodynamics states that the entropy of all substances inpure crystalline state at 0 K is zero.

6. Free energy the amount of energy involved in a process that is available to dowork.PREVIEWING ACTIVITIES

A. Let the students to list down processes or reactions happening around them thatare:

1. spontaneous.2. non-spontaneous.

B. Discuss the Second Law of Thermodynamics.C. Discuss the Third Law of thermodynamics.D. Pose the Guide Questions that the students will answer after viewing the

episode. Ask them to focus on finding the answers to the guide questions as they

watch the video.


1. What is entropy?Entropy is the thermodynamic quantity that refers to the degree of

randomness or disorder of a system.

2. Is entropy a measurable quantity?

The measurable quantity is the change in entropy or the change in the

disorder or randomness of the system when a certain process occurs.

However, there are reported values of absolute entropy of substances, whose

determination is made possible by setting as a reference point when entropyof all substances is set at zero and this is when substances are in pure

crystalline states at absolute zero, according to the Third Law of

Thermodynamics.

3. What makes a reaction spontaneous?

A reaction or process is spontaneous if there is a net free energy released

when it occurs. This is indicated by a negative value for G in the Gibbs

Free Energy equation.


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VIEWING ACTIVITIES

Let the students view the segments on

Entropy and the Second and Third Laws of Thermodynamics at2:30 - 10:30 and 10:40 - 23:20.

POST-VIEWING ACTIVITIES

A. Discuss the answers to the Guide Questions.

B. Discuss how the Second Law of Thermodynamics is applied as per theviewing activity.

C. Discuss how the Third Law of Thermodynamics is applied as per the videoclip.

ASSESSMENT

Quiz. Choose the letter corresponding to the best answer.

For Questions 1-3, consider the hypothetical endothermic reaction:

4A(g) + B2(g) 2A2B(g)

1. The change in entropy of the system is most likelyA. negative. C. zero.B.positive. D. cannot be determined.

2. The reaction is

A. spontaneous at all temperatures. C. spontaneous only at high

temperatures.B. non-spontaneous at any temperature. D. non-spontaneous at low

temperatures.

3. The reverse of the above reaction is therefore

A. spontaneous and exothermic. C. spontaneous and endothermic.

B. non-spontaneous and exothermic. D. non-spontaneous and endothermic.

4. Which of the following results in a decrease in entropy?A. Evaporation of water C. Melting of iceB. Freezing of water D. Sublimation of dry ice


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5. At what temperature would you expect the melting of ice to be non-spontaneousat normal atmospheric pressure?

A. -10oC C. 10oCB. 1oC D. 100oC

6. Which of the following processes is non-spontaneous at 25oC?A. Condensation of water vapor C. Burning of candleB. Ripening of mango D. Rusting of iron

7. An increase in entropy occurs whenA. solids form from the freezing of pure liquids.B. the temperature of a substance decreases.C. the number of molecules of gas present in the system increases as a result of

a chemical reaction.

D. a gas condenses into a liquid.8. Which of the following is true about free energy?

A. It is available or unused energy. C. It is released by the system.B. It is absorbed by the system. D. It is zero at absolute zero

temperature.

9. Which of the following statements is false?A. A reaction that is endothermic and involves an increase in entropy can be

spontaneous at high temperatures.B. In increase in molecular motion results in an increase in entropy.C. Spontaneity is favored by a decrease in both entropy and enthalpy.D. The entropy of the universe is increased when spontaneous reactions occur.

10.At absolute zero temperature,A. most substances have solidified. C. particles are no longer moving.B. melting of ice is spontaneous. D. the free energy of all processes is

zero.

ANSWER KEY

Quiz.

1. A 2. B 3. A 4. B 5. B

6. A 7. C 8. A 9. C 10. C


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REFERENCES

Brown, T. L., LeMay, H. E. & B. E. Bursten. (2004). Chemistry: the central science.

NJ: Prentice Hall, Inc.

Silberberg, M. S., (2000). Chemistry: the molecular nature of matter and change.

NY: McGraw-Hill, Inc.

Useful Website

Volland, W. Online introductory chemistry. http://scidiv.bcc.ctc.edu/wv/7/0007-

003-free _energy.htm


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Chapter 6: Rates and L imits of Change

EPISODE 20: RATES OF REACTION

OVERVIEW

Chemical reactions are processes in which one or more substances are converted

into other substances. Evidences that indicate the occurrence of chemical changesinclude change in color, evolution of gas, formation of precipitate or change in

temperature of the reacting system. Some of these evidences can be observedimmediately after the reactants come together for certain reactions, but for other

reactions, the observables are not readily apparent because chemical reactions may

occur in a snap of a finger or may require many years to reach completion.

This episode introduces us to chemical kinetics, the study of the rates andmechanisms of chemical reactions. The factors affecting rates of reactions are

investigated and explained using the Collision Theory.

OBJECTIVES

At the end of this lesson, the student should be able to:1. define rate of reaction;2. identify parameters used to measure rate of reaction;3. enumerate and explain the factors affecting the rates of chemical reactions;4. state the Collision Theory of molecules;5. use the Collision Theory to explain the factors affecting reaction rates; and6. cite applications of chemical kinetics in everyday processes.


This video lesson makes use of concepts learned in Episode 10 - Indicators of

Chemical Change, 16 - Energy Conservation, and 17 - Energy and Change.

SCIENCE PROCESSES

Using space-time relationships

InferringClassifying

Interpreting data

Predicting

Observing

ExperimentingMeasuring

Controlling variables



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change that happens in a given time interval. The rate of a chemical reaction is

measured in either of two ways: by the decrease in the amount or concentration of a

reactant or by the increase in the amount of a product within a given unit of time.

Consider the hypothetical chemical reaction:A B

(reactant) (product)Rate of reaction = change in the concentration of A per unit interval of time or

Rate of reaction = change in the concentration of B per unit interval of time

We usually look for a property of a reactant or a product of the reaction that can be

monitored qualitatively or quantitatively through the use of some measuring

instrument. For example, if a reaction produces a colored product, you may follow

the progress of a reaction by monitoring the increase in intensity of the color of the

reacting mixture as more product is formed. If the reactant is colored, but none of

the products are, the rate of disappearance of color can be used to measure rate.

Some reactions produce gaseous products and the rate of gas evolution can befollowed.

An example of a qualitative determination of how fast a reaction occurs is

demonstrated in the video lesson. In the experiment shown, the burning of the wood

splinter in air is compared with the burning of another wood splinter inside a test

tube containing pure oxygen. The flame produced is much brighter and the burning

of the wood is quicker inside the test tube with pure oxygen. It can be inferred fromthese observations that the rate of reaction inside the test tube is greater than the oneoccurring in open air.

A more quantitative

comparison of rates can be

done by following the

changes in the concentration

of the reactants or products in

the reaction and plotting thedata in the form of decay or

growth curves of substances.A decay curve of the reactant

indicates the decrease inamount of reactants still

remaining as the reaction

proceeds as illustrated in the

figure at the right.


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Factors Affecting Reaction Rates. The rate of a reaction is affected by thefollowing factors: concentration of reactants, temperature, and the presence of

catalyst.

Concentration of Reactants. The burning of the wood splinter shown in the videolesson is a good demonstration of the effect of concentration of reactant, in this case

oxygen, on the rate of the reaction. Burning, or in more technical terms, combustion,is the reaction of substances with oxygen. The flame produced from the burning of

the wood splinter in open air, which has only about 20% oxygen, is less bright andthe burning is slower compared to that in the tube containing pure oxygen.

A second laboratory demonstration on the effect of concentration of reactants that is

shown in the video lesson is the reaction between hydrochloric acid, HCl in solution,

and zinc, Zn, metal. The equation of the reaction is:

HCl(aq) + Zn(s) Zn2+

(aq) + 2Cl-(aq) + H2(g)

Since one of the products of this reaction is a gas, H2, the determination of the rate

of reaction is done by getting the volume of gas produced for a certain period oftime. An improvised gas buret is used to measure the volume of the gas formed.

In the experiment, a constant amount of Zn is dropped in each of the test tubes

containing different concentrations of HCl. The rates of reaction at different

concentrations of HCl were compared by measuring the amount of hydrogen gas

produced in each test tube for the same length of reaction time. The results of theexperiment showed that more hydrogen gas was produced from the reaction mixture

with a higher concentration of HCl. Thus, the higher the concentration of the HCl

solution, the faster is the chemical reaction.

As the concentr ation of r eactants incr ease, the rate of the reaction al so incr eases.

Temperature. The same experimental set-up was used in the video lesson todemonstrate the the effect of temperature on the rate of reaction of HCl and Zn.

Three different flasks, each containing the same volume of HCl solution and Zn aresubjected to different temperatures: hot water bath (40oC), room temperature (30

oC),

and ice water bath (10oC). The results showed that the rate of gas evolution is

greater at higher temperature.

One very familiar application of the effect of temperature on rates of reaction isrefrigerating perishable foods such as milk. The bacterial reactions that lead to the

spoiling of milk proceed more rapidly at room temperature than at the lower

temperature of a refrigerator.

The higher the temperature, the faster is the reaction.


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Catalyst . The presence of a catalyst is another factor that influences the rate of achemical reaction. A catalyst is a substance that speeds up a chemical reaction, but is

not consumed in the reaction.

Hydrogen peroxide or agua oxigenada undergoes decomposition by itself, especiallywhen exposed to bright light. The decomposition reaction is accompanied by the

formation of bubbles as oxygen gas is formed and escapes from the liquid.

(slow reaction) H2O2(l) 2H2O(l) + O2(g)

In a dark bottle, the reaction is very slow. But if manganese dioxide is added,vigorous bubbling is observed.

MnO2

(fast reaction) H2O2(l) 2H2O(l) + O2(g)Manganese dioxide is not an added reactant to the process but a catalyst and is not

consumed in the reaction. An evidence of this is that one can recover the substance

from the system. There is no change in its amount after being recovered from thereaction vessel.

A catal yst can i ncrease the rate of reactionby providing an alternative path for the reactant molecules.

This new path is characterized by a lower activation energy requirement. The

activation energy is the minimum energy that reactant molecules must acquire toenable them to undergo bond breaking and electron rearrangements during thereaction. Unless a reactant molecule has enough energy to meet this minimum, it

cannot be converted into the product. When a catalyst opens an alternative route that

is of lower activation energy for the reactants, then more reactant molecules canmeet the energy requirement and react to form the products.

TheColl ision Theory of M olecules. The effects of the aforementioned factors onthe rate of reactions are bestexplained by the Coll ision Theory. The theory suggeststhat for a chemical reaction to occur, the reactant molecules must collide. The

collision must be effective so that a successful reaction results. How can collisionsbecome effective? The colliding molecules must be properly oriented and that the

collisions have sufficient energy to meet the activation energy requirement.

According to the Collision Theory, an increase in concentration of reactants resultsin an increase in reaction rates because the increase in the number of particles in the

reaction mixture results in an increase in the frequency of collisions. An increase inthe temperature of the reaction increases reaction rate because the reacting particles

would have greater kinetic energy and thus would collide with one another more

frequently and collide with greater impact. When colliding particles have moreenergy, the probability of collisions becoming effective is also greater. Catalysts are


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PRE-VIEWING ACTIVITIES

A. Initiate a discussion about the spectacular display of colors in fireworks. Let thestudents tell their own stories.

B. Present pictures of different chemical reactions. Lead the students to note thatthese chemical reactions differ in terms of rate.

C. Pose the Guide Questions that the students will answer after viewing theepisode. Ask them to focus on finding the answers to the Guide Questions as

they watch the video clip.


1. What is rate? How is the rate of a moving object measured?Rate is a measure of how much change has occurred per unit time. For amoving object, rate expresses how much an object has been moved or haschanged position over time. It is measured by getting the distance it hastraveled divided by the time.

2. What is the rate of chemical reaction? How is it determined?The rate of a chemical reaction refers to the speed by which a chemicalreaction proceeds. Rate of reaction is determined by the decrease in the

concentration of a reactant or the increase in the concentration of a productdur ing a given unit of ti me.

3. Explain how the decay and growth curves may be useful in describing therate of reaction.

The decay and growth curves of substances are useful in quantitativelydescribing changes in the concentrations of the reactants and products in areaction. The decay curve indicates how much of the reactants are used upwhile the growth curve indicates the increase in the amount of the productsformed.

4. Name the factors that affect reaction rate.The main factors that affect reaction rate are concentration of substances,temperatur e, and the presence of a catalyst.

5. How do these factors affect the rate of a chemical reaction?In general, an increase in the concentration of the reactants increases therate of reaction. As temperature incr eases, the rate of chemical reaction alsoincreases. A catal yst speeds up a r eaction.

6. What is a catalyst? Why cant it be considered as another reactant in achemical reaction?


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A catalyst is a substance that is present in and speeds up a chemicalreaction. It cannot be considered as another reactant in a chemical reactionbecause normally, a reactant is consumed, slowly disappearing in thereaction while a catalyst is not consumed. It remains unchanged in the

course of the reaction.

7. What is explained by the Collision Theory?The Col l ision Theory states that for a chemical reaction to occur, there mustbe effective collisions among the reacting molecules. Effective collisionmeans that the molecules are properly oriented during collision and thatthese molecules must possess the minimum energy of acti vation.

8. How can you relate the Collision Theory to the factors controlling reactionrate?

The Collision Theory explains how an increase in concentration of thereacting specie and the reaction temperature can affect reaction rate. Withincreased concentration of reactant molecules, the chances for molecularcollision consequently increase. The more collisions occurring, the greaterare the chances for effective col l isions, re

Chem Volume 2

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