THERMODYNAMICS - umlub.pl · 11/14/2018 2 Definitions 5 Thermochemistry deals with changes in heat...

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THERMODYNAMICSThermochemistry

OBJECT OF THE THERMODYNAMICS

In thermodynamics we study the energy changes that accompany physical and chemical processes.

Usually these energy changes involve heat—hence the “thermo-” part of the term.

There are the two main aspects of thermodynamics.

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THERMOCHEMISTRY

The first aspect is thermochemistry.

This practical subject is concerned with how we observe, measure, and predict energy changes for both physical changes and chemical reactions.

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FUNDAMMENTAL ASPECT OF THERMODYNAMICS

• The second aspect is addressed to a more fundamental aspect of thermodynamics.

• How to use energy changes to tell us

whether or not a given process can occurunder specified conditions

• to give predominantly products (or reactants)

• how to make a process more (or less) favourable.

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Definitions5

Thermochemistry deals with changes in heat during

chemical reactions.

A main goal of the thermochemistry study is to determine

the quantity of heat exchanged between a system and its

surroundings.

The system is the part of the universe being studied,

while the surroundings are the rest of the universe that

interacts with the system.

System and surroundingsWM_ THERMODYNAMICS 6

System and surroundings7

Open system

An open system is a system that freely exchanges energy and matter with its surroundings.

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WM_ THERMODYNAMICS 9

Open systems

Closed system

A closed system exchanges energy but not matter with an outside system(surroundings). Although it is typically portion of larger system, it is not in complete contact with it.



Isolated systemWM_ THERMODYNAMICS 12

An isolated system can exchange neither energy

nor matter with its surroundings (an outside

system). While it may be portion of larger system,

it does not communicate with the outside in any

way.

Examples of such system type are:

physical universe and a closed thermos bottle

(though its isolation is not perfect).

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13Comparison of systems

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Quiz15

•Q,

•A closed system contains 2g of ice. Another 2g

of ice are added to the system. What is the final

mass of the system?

•Q

•An isolated system has an initial temperature of 30oC.

It is then placed on top of a bunsen burner for an hour.

What is the final temperature?


System description System type

Coffee in perfectly closed Thermos® flask

Combustion of gasoline in car engine

Mercury in thermometer

Living plant

Electric battery

Q. Identify system type (open, closed or isolated) from

description below and fill the empty space in the table

below

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Q. Which type of thermodynamic system is:

1. an ocean?

2. an aquarium?

3. a pizza delivery bag?

4. a greenhouse?

5. a man ?

ENERGYMatter and energy

18WM_ THERMODYNAMICS

MATTER AND ENERGY

Matter is anything that has a mass and occupies some space. All bodies consist of a matter.

Mass is a measure of the quantity of a matter in a sample of any material.

The more massive an object is, the more force is required to put it in a motion.

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Energy

Energy is measure of the ability of a body or system to do work or produce anychange. No activity is possible without energy.

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ENERGY FORMS

Energy can take many forms:

electrical energy,

radiant energy (light),

nuclear energy,

chemical energy.

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KINETIC ENERGY

Commonly we classify energy into two general types: kinetic and potential.

Kinetic energy is the energy of motion. The kinetic energy of an object is equal to one half its mass, m, times the square of its velocity, v.

The heavier a hammer is and the more rapidly it moves, the greater its kinetic energy and the more work it can accomplish.

2

2

1vmEk

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Kinetics energy in questions

Q.

The kinetics energy of solid body with the mass of 5 kg which moved with speed 8 m s-1 is equal:

a) 40 kg m s-1; b) 320 J; c) 160 J

d) 160 kg m2s-2


POTENTIAL ENERGY

Potential energy (EP) is the energy thata system possesses by virtue of its position or composition. It is storedenergy.

The work that we do to lift an object is stored in the object as energy.

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Ep= m g h

Where: m – mass (kg); h – body movement (change of

height, m); g – gravitational acceleration, (10 m s-2 )

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Potential energy in question

Q. When a bucket with 10 kg of water is picked up at the height of 1 m the potential energy is as follows:

a) 10 kg m; b) 100 J; c) 100 kg m2s-2;

d) 100 kg

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EXAMPLE: EpEk

If we drop a hammer, its potential energy is converted into kinetic energy as it falls, and it could do work on something it hits—for example, drive a nail or break a piece of glass.


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Q. What type of energy does a stationary

pencil contain? falling pencil?

Chemical changes always involve energy changes. However, some energytransformations do not involve chemical changes at all.

For example, heat energy may beconverted into electrical energy or into mechanical energy without any simultaneous chemical changes.

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The Law of Conservation of Matter and Energy

With the dawn of the nuclear age in the 1940s, scientists, and then the world, became aware that matter can be converted into energy.

In nuclear reactions, matter is transformed into energy.

The relationship between matter and energy is given by Albert Einstein’s now famousequation:

E = m c2

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Law of Conservation of Energy

Many experiments have demonstrated that all of the energy involved in any chemical or physical change appears in some form after the change.

These observations are summarized in the Law of Conservation of Energy:

Energy cannot be created or destroyed in a chemical reaction or in a physical change.

It can only be converted from one form of energy to another.

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The Law of Conservation of Matter and Energy

Now that the equivalence of matter and energy is recognized, the Law of Conservation of Matter and Energy can be stated in a single sentence:

The combined amount of matter and energy in the universe is fixed.


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MEANING OF THE ENERGY IN THE LIFE

Energy is very important in every aspect of our daily lives.

The food which we eat supplies the energy to sustain life with all of its activities and concerns.

The availability of relatively inexpensive energy is an important factor in our technological society.

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„ACCOUNTING OF ENERGY’

• The concept of energy is at every heart of science.

• All physical and chemical processes are accompanied by the transfer of energy.

• Energy cannot be created or destroyed.

• We must understand how to do the “accounting” of energy transfers from one body or one substance to another or from one form of energy to another.

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ENERGY - DEFINITION

We can define energy as follows:

Energy is the capacity to do work or to transfer heat.

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Heat, energy, work units

The amount of heat transferred in a process is usually expressed in joules or in calories.

The SI unit of energy and work is the joule (J), which is defined as 1 kg m2/s2.

Not SI unit of energy e.g. 1 cal = 4.184 J


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Heat as a form of energyWM_ THERMODYNAMICS 38

Heat is a form of energy that always

flows spontaneously from a hotter

body to a colder body—never in the

reverse direction

Heat transfer

Heat transfer concerns the generation, use, conversion, and exchange of heat(thermal energy) between physical systems.

Heat transfer is classified into various mechanisms, such as thermal conduction(diffusion), thermal convection, thermalradiation and by phase changes.


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Conduction is the transfer of thermal energy through direct contact between particles of a substance, without moving the particlesto a new location. •Convection is the transfer of thermal energy through movement of particles from one location to another

•Radiation is the emission of energy as waves or particles or rays. TEMPERATURE vs. HEAT

Temperature measures the intensity of a heat, the “hotness” or “coldness” of a body.

A piece of metal at 100°C feels hot to the touch, whereas an ice cube at 0°C feels cold.

Why? Because the temperature of the metal is higher, and that of the ice cube lower, than body temperature.


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State functionsWM_ THERMODYNAMICS 42

It turns out that the energy of an object depends only on the object’s current condition.

The complete list of properties that specify an object’s current condition is known as the state of the object. In chemistry it is usually enough to specify the object’s pressure, temperature, volume, and chemical compositions (numbers of moles) to give the state of the object.


STATE FUNCTION

Any property of a system that depends only on the values of its state functions is also a state function.

For instance, the volume of a given sample of water depends only ontemperature, pressure, and physical state; volume is a state function.

We shall encounter other thermodynamic state functions.


THERMODYNAMIC STATE

The thermodynamic state of a system is defined by a set of conditions that completely specifies all the properties of the system.


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THERMODYNAMIC STATE

This set commonly includes:

temperature, T,

pressure, P,

volume, V,

composition (identity and number of

moles of each component), n,

physical state (gas, liquid, or solid) of

each part of the system.

46WM_ THERMODYNAMICS WM_ THERMODYNAMICS 47

Ice liquid water

Steam

Changes in physical

state of water due

to temperature

changes

STATE FUNCTIONS

The properties of a system—such as P, V, T—are called state functions.

The value of a state function depends only on the state of the system and not on the way in which the system came to be in that state.

A change in a state function describes a difference between the two states.

48 49

1

2

3

Direct

Indirect

Indirect

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STATE FUNCTIONS

For instance, consider a sample of one mole of pure liquid water at 30°C and 1 atm of pressure.

If at some later time the temperature of the sample is 22°C at the same pressure, then it is in a different thermodynamic state.

Thus change in temperature is equal to:

∆𝑡 = 𝑡𝑓𝑖𝑛𝑎𝑙 - 𝑡𝑖𝑛𝑖𝑡𝑖𝑎𝑙

We can tell that the net temperature change is 8°C.


STATE FUNCTION

It does not matter whether:

(1) the cooling took place directly (either slowly or rapidly) from 30°C to 22°C,

or (2) the sample was first heated to 36°C, thencooled to 10°C, and finally warmed to 30°C, then cooled to 22oC

or (3) any other conceivable path was followedfrom the initial state to the final state.


STATE FUNCTION

The change in other properties (e.g., the pressure) of the sample is likewise independent of path.


STATE FUNCTIONS

The most important use of state functions in thermodynamics is to describe changes.

We describe the difference in any quantity, X, as

When X increases, the final value is greater than the initial value, so ΔX is positive;

a decrease in X makes ΔX a negative value.

INITIALFINAL XXX


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ThermochemistryWM_ THERMODYNAMICS 54

CALORIMETRY, thermochemistry

We can determine the energy change associated with a chemical or physical process by using an experimental technique called calorimetry.

This technique is based on observing the temperature change when a system absorbs or releases energy in the form of heat. This is in turn the effect of chemical or physical process under study.


The experiment is carried out in a device called a calorimeter, in which we measure the temperature change of a known amount of substance (often water) which specific heat is known.


Specific heat

The specific heat, c, is the amount of heat per unit of mass required to raise the temperature by one degree Celsius orKelvin with no change in phase.


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SPECIFIC HEAT [c]


SPECIFIC HEAT

The specific heat of each substance, a physical property, is different for the solid, liquid, and gaseous phases of the substance.

For example, the specific heat of:

ice is 2.09 J/g °C near 0°C;

liquid water is 4.18 J/g °C;

steam is 2.03 J/g °C near 100°C.


WM_ THERMODYNAMICS 61SPECIFIC HEAT

Substance c in J/g KMolar CJ/mol K

Aluminum 0.900 24.3

Copper 0.386 24.5

Gold 0.126 25.6

Lead 0.128 26.4

Silver 0.233 24.9

Zinc 0.387 25.2

Mercury 0.140 28.3

Alcohol(ethyl) 2.400 111.0

Water 4.186 75.2

Ice (-10 C) 2.050 36.9

Specific heats and molar heat capacities for

various substances at 293 K (20oC)


Q. How many heat is needed to heat up 10 g of liquid

water from 10oC to 40oC.

[c of liquid water is 4.18 J g-1 oC-1].

a) 418 J; b) 1254 J; c) 1672 J

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Answer:

Q = m c ΔT

Q = (10 g) . (4.18 J g-1 oC-1) (40 – 10)oC

Q = 10 . 4.18 . 30 = 1254 J

Specific heat

The specific heat, c, is the amount of heat per unit of mass required to raise the temperature by one degree Celsius orKelvin with no change in phase.


Changes in phase

Changes in phase (physical state) absorb or liberate relatively large amounts ofenergy (see Figure –NEXT SLIDE).

65WM_ THERMODYNAMICS 66

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THERMAL EQUILIBRIUM

A hot object, such as a

heated piece of metal (a), is

placed into cooler water.

Heat is transferred from the hotter

metal bar to the cooler water until

the two reach the same

temperature (b).

We say that they are then at thermal equilibrium


A 34 gram piece of an unknown metal

absorbs 351.56 Joules of energy when

the temperature increased from 10oC to

32oC. What is the specific heat of the

substance?

Hint: You are solving for Specific Heat

(Cp) not heat absorbed.

Quiz


Q.

A 385 grams chunk of iron is heated to

97.5oC. Then it is immersed in 247 gram of

water originally at 20.7oC. When thermal

equilibrium has been reached, the water and

iron are both at 31.6oC. Calculate the specific

heat of iron.


Solution:

The number of heat gained by water from

temperature 20.7oC to 31.6o C =the amount of heat

which is lost by the iron.

Qwater = (247 g). (4.18 J g-1 oC-1) (31.6-20.7oC)

Qwater = (247) . (4.18). (10.9) = 11253.8 J

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If a 25.2 g piece of silver absorbs 365 J of heat,

what will be the final temperature

of the silver if the initial temperature is 22.2 oC?

The specific heat of silver is 0.235 J/g K


Expansion work.

Pressure-volume work of gas


Atmospheric pressure

pin

Gas under

pressure

Atmospheric pressure

Gas at atmospheric pressure

1 2

Heat capacity of calorimeter = 8.101 kJ/oC, mesured in separated experiment

Constant pressure


Run Pin position Initial buckettemperature(oC)

Finalbuckettemperature(oC)

1 locked 24.00 28.91

2 unlocked 27.32 31.54

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q for run 1

q1= C Δt = 8.101 kJ/oC x (28.91 – 24.00)oC = 39.8 kJ

heat is released by reaction thus q= -q = -39.8 kJ

q for run 2 at constant pressure

q2= C Δt = 8.101 kJ/oC x (31.54 – 27.32)oC = 34.2 kJ

heat is released by reaction thus q= -q = -34.2 kJ

Missing 5.6 kJ ????

Expansion work (pressure –volume work)

w= -PΔV


ΔE = q + w

Work and heat are simply alternative ways

to transfer energy.

First law of thermodynamics = law of

conservation of energy

FIRST LAW OF THERMODYNAMICS

Some important ideas about energy are summarized in the First Law of Thermodynamics.

Energy is neither created nor destroyed in ordinary chemical reactions and physicalchanges.


THE UNIVERSE, SYSTEM, SURROUNDINGS

The substances involved in the chemical and physical changes that we are studying are called the system.

Everything in the system’s environment constitutes its surroundings.

The universe is the system plus its surroundings.


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FIRST LAW OF THERMODYNAMICS

The system may be thought of as the part of the universe under investigation.

The First Law of Thermodynamics tells us that energy is neither created nor destroyed.

Energy is only transferred between the system and its surroundings.

79WM_ THERMODYNAMICS WM_ THERMODYNAMICS 80

Q. The first law of thermodynamics

states that energy is

a. increased during any process

b. decreased during any process

c. conserved during any process


FUNDAMMENTAL ASPECT OF THERMODYNAMICS

Potential energy of atom

An electron in an atom has potential energy because of the electrostatic force on it that is due to the positively charged nucleus and the other electrons in that atom and surrounding atoms.

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ATOMIC LEVEL of energy

The atomic or molecular level, we can think of each of these as either kinetic or potential energy.

The chemical energy in a fuel or food comes from potential energy stored in atoms due to their arrangements in the molecules.

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INTENSITY OF HEAT

Many forms of energy can be interconverted

and that in chemical processes, chemical

energy is converted to heat energy or vice

versa.

The amount of a heat a process uses

(endothermic) or gives off (exothermic)

can tell us a great deal about that process.

For this reason it is important for us to

be able to measure the intensity of the

heat.

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https://www.quora.com/What-is-difference-between-endothermic-and-

exothermic-reaction-if-both-require-activation-energy

ACTIVATION ENERGY

In such reactions, the total energy of the products is lower (for exothermic) orgreater (endothermic) than that of the reactants by the amount of energy as a heat released or absorbed.

Some initial activation (e.g., by heat)is needed to get these reactions started. This amount of energy is called activation energy.


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THERMOTERMIC REACTION

The amount of heat shown in thermochemical equation always refers to the reaction for the number of moles of reactants and products specified by thecoefficients.


Thermochemical equation

𝑁2(𝑔) + 3 𝐻2(𝑔) → 2𝑁𝐻3(𝑔) Δ𝐻0 = -92.38 kJ

2𝑁2(𝑔) + 6 𝐻2(𝑔) → 4𝑁𝐻3(𝑔)Δ𝐻0 = 2 x(-92.38 kJ ) = -184.8 kJ

1/2𝑁2(𝑔) +3

2𝐻2 𝑔 → 𝑁𝐻3 𝑔

Δ𝐻0 = (0.5) -92.38 kJ = -46.19 kJ/ mole


2H2(g) + O2(g)


2H2O(l)

ΔH0 = -571.8 kJ

H2 g + ½ O2(g)→H2O(l)

ΔH0 = ? kJ/mole

EXOTHERMIC REACTIONS

Reactions that release energy in the form of heat are called exothermic reactions.

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CH4(g) + 2O2(g)→ CO2(g) + 2H2O(l) + 890 kJ/mol

reagents products

exothermic reaction at constant pressure.

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Exothermic reaction Endothermic reaction

Reactions that absorb energy in the form of heat are called endothermicreactions.


ENDOTHERMIC REACTION

NH4NO3(s) + 26 kJ NH4NO3(aq)


Reagents product

H2O

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Exothermic Reaction

Endothermic Reaction

Energy absorbed orreleased

Energy is released.It is a productof the reaction.Reaction vessel becomes warmer.

Temperature inside reaction vessel increases.

Energy is absorbed.It is a reactant of the reaction.Reaction vessel becomes cooler.

Temperature inside reaction vessel decreases.

Relative Energy of reactants & products

Energy of the reactants is greater than the energy ofthe productsH(reactants) > H(products)

Energy of the reactants is less than the energy of the productsH(reactants) < H(products)

Sign of HH = H(products) -H(reactants)

= negative (-ve)

H = H(products) -H(reactants)

= positive (+ve)

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Reversible reactionWM_ THERMODYNAMICS 95

N2 + 3 H2 ↔ 2NH3

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Exothermic Endothermic

Energy of the reactants

(N2 & H2) is greater

than the energy of the

products (NH3). Energy

is released.

Energy of the reactants

(NH3) is less than the

energy of the products

(N2 and H2). Energy is

absorbed.

En

erg

y p

rofile

s

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Exothermic processes Endothermic processes

making ice cubes melting ice cubes

formation of snow in clouds conversion of frost to water vapour

condensation of rain from water

vapourevaporation of water

mixing sodium sulfite and bleach baking bread

rusting iron cooking an egg

burning sugar producing sugar by photosynthesis

mixing water and strong acids mixing water and ammonium

nitrate

mixing water with an anhydrous salt making an anhydrous salt from a

hydrate

crystallizing liquid salts (as in

sodium acetate in chemical

handwarmers)

melting solid salts

Author: Fred Senese [email protected];

http://antoine.frostburg.edu/chem/senese/101/thermo/faq/exothermic-endothermic-examples.shtmlWM_ THERMODYNAMICS 98

Q. Chemical reactions that absorb heat energy are called __________ . a. exothermicb. eltothermicc. endothermic

Q. Electrolysis requires energy to make it work. This means it is...a) an endothermic reaction b) an exothermic reaction c) an eltothermic reaction d) a chemical reaction

http://antoine.frostburg.edu/chem/senese/index.shtml

mailto:[email protected]

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Q. Which of the following is an endothermic

reaction?

a) Burning propane in a gas grill

b) Photosynthesis

c) Baking bread

d) Cooking an egg

e) Electrolysis of water

Q. A. What is the change energy of the

sausage after heating, if original energy is

4 kJ and 20 kJ is added to it?

B. What is the total energy content of

sauage after heating?

a) 16 kJ; b) 4 kJ; c) 20 kJ; d) 24 kJ


ENTHALPY

Most chemical reactions and physical

changes occur at constant (usually

atmospheric) pressure.

The quantity of heat transferred into or out

of a system as it undergoes a chemical or

physical change at constant pressure, qp,

is defined as the enthalpy change, H, of

the process.


ENTHALPY - HEAT

The enthalpy change is equal to the enthalpy or “heat content,” H, of the substances produced minus the enthalpy of the substances consumed.


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Difference between energy and enthalpy change

∆𝐸 − ∆𝐻 = −𝑝∆𝑉


Pressure volume

work

The only time when ΔH and ΔE differs by a significant amount is when gases are formedor consumed in a reaction. In such situation we applied ideal gas law and obtain formula for ΔH as follows:

∆𝐻 = ∆𝐸 + ∆𝑛𝑔𝑎𝑠RT

EHTHALPY AS STATE FUNCTION

It is impossible to know the absolute enthalpy (heat content) of a system.

Enthalpy is a state function, however, and it is the change in enthalpy in which we are interested. This can be measured for many processes.


CHANGE OF ENTHALPY109WM_ THERMODYNAMICS WM_ THERMODYNAMICS 110

Calculate the enthalpy of the ammonium nitrate

decomposition. The reaction is

and the enthalpies of the three compounds are given in

Table 1.

ΔH = [ 82+ 2(-242)] – [-366] = -36 kJ

N2O H2O NH4NO3

products reactant

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111

If you reverse the previous reaction,

the sign of the enthalpy of the reaction is reversed:

Δ H = +36 kJ


http://www.cliffsnotes.com/sciences/chemistry/chemistry/thermodynamics/e

nthalpy


Po

ten

tiale

ne

rgy

reactants

products

ΔH > 0

products

Po

ten

tiale

ne

rgy

reactants

ΔH < 0

Reaction progress Reaction progress

EndothermicExothermic

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There are two ways of looking at what happens

to the enthalpy:

If the reaction is exothermic the products

have minimum enthalpy and the formation of

products (move toward the right) is favourable

If the reaction is endothermic the reactants

have minimum enthalpy and the formation of

products (move toward the right) is

unfavourable.

In this case the formation of reactants (move

toward the left) is favourable.

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•Calculate the enthalpy change for the following

reaction and classify it as exothermic or

endothermic.


Compound ΔH0

MgCl2 (S) -642 kJ/mol

H2O (l) -286 kJ/mol

MgO (S) -602 kJ/mol

HCl (g) -92 kJ/mol

Standard enthalpies of formation

ΔH = [ΔH0 (MgO) + 2 ΔH0 (HCl)] – [ ΔH0 (MgCl2) + ΔH0(H2O)]

End of part 1


THERMODYNAMICS - umlub.pl · 11/14/2018 2 Definitions 5 Thermochemistry deals with changes in heat...

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