thermodynamics

THERMODYNAMICS

SPECIFIC LEARNING OBJECTIVEAt the end of the session the student should be able to explain:•Energy•The first law of thermodynamics•Entropy•Free energy

Thermodynamic

In ChemistrySystem

In Living

System

THERMODYNAMICS

Thermodynamic is the law that formulated from observation on

conversion of energy from one form to the other. i.e. transduction

What is the energy?

1. Energy is a much used term, but it represents a rather abstract

concept.2. Energy is usually defined as the capacity to do work.3. Chemist define work as directed energy change resulting from a process

The type of Energy

1. Kinetic energy

2. Radiant energy

3. Thermal energy

4. Chemical energy

5. Potential energy

Definition of type energy

1. Kinetic energy – the energy produced by a moving object2. Radiant energy : comes from the

sun (solar energy) and is Earth’s primary energy source. Solar energy heats the atmosphere and Earth’s surface, stimulates the growth of vegetation through the process known as photosynthesis, and influences global climate patterns.

continued

The Activated Complex form the reaction: A + B AB# P

Definition of type energy

3.Thermal energy is the energy associated with the random motion of atoms and molecules.

4. Chemical energy is stored within the structural units of chemical

substances; its quantity is determined by the type and arrangement of atoms in the substance being considered.

5. Potential energy is energy that is also available by virtue of an object’s position.

Conclusion of Energy

• All forms of energy can be interconverted (at least in principle) from one form to another

Scientists have concluded that energy can be neither destroyed nor created.

• Thermodynamic Law

THERMODYNAMICS

Thermodynamic I

Thermodynamic II

Thermodynamic III

Thermodynamic I is the law of conservation of energy

Work Heat

Work and heat are not state functions

• The relationship between chemical energy and other forms of energy, with examples.

Energy change in chemical reactions

Almost all chemical reactions absorb or produce (release) energy, generallyIn the form of heat.• Heat is the transfer of thermal energy between two bodies that are at differenttemperatures. Although ”Heat” itself implies thetransfer of energy, we customarily talk of ”heat absorbed” or ”heat released” when describing the energy changes that occurduring a process.

Energy changes associated with chemical reactions

System Surroundings

SYSTEM AND SURROUNDING

System we mean that the part of the world we are investigating.

Surrounding we mean everything else

system

Surrounding

Open Close

be

Three type of systems

Isolated

System be open

• Two of this examples are the examples of open system:

1.e.g. in the living organism, which takes up nutrients, releases the waste products, and generates work and heat.

2.An example in body, the body takes up nutrient, and then release urine which contains toxin, carbon dioxide, and so on.

System be closed

Example of close system:• An example of close system is living

of an microorganism, it was sealed inside a perfectly insulated box, it will, together with the box, constitute a closed system.

q, to be the manner of energy transfer that results from a temperature difference between the system and its surrounding

HEAT

Positive and negativesign of heat

Heat input to a system is considered a positive quantity

Heat evolved by a system isconsidered a negative quantity.

W O R K

w, to be the transfer of energy between the system of interest and its surroundings asa result of existence of unbalanced forcesbetween two.

Positive and negative sign of work

if the energy of the system isdecreased by the work, or thesystem does work on thesurroundings, or that work isdone by the system, and wetake it to be a negativequantity

If the energy of the system is increased by the work, we say that work is done on the system by surroundings, and we take it to be a positive quantity

• The effect of work is equivalent to the raising or lowering of mass in

the surroundings.

• Work is done by the system because the mass is raised

work is done on the system because the mass is lowered.

E N E R G Y

Thermodynamic I study of conservation of energy

Energy is a state functionIt is a property that depends only upon the state of the system, and not upon how the system was brought to that state, or upon

the history of the system.

The first law of thermodynamic

•∆ U = q + wwhich is essentially a statement of the

law of conservation of energy.

1. The term ∆ U represents the change of internal energy of the system, 2. q is the thermal energy (heat) added to the system, and w is the work done on the system.

Where :

The chemical reactions that need energy

1. The Photoelectric Effect this is mystery in physics.

Experiments had already demonstrated that electrons were ejected from the surface of certain metals exposed to light of at least a certain minimum frequency.

Einstein suggested that a beam of light is a stream of particles. These particles of light are called photons. Using Planck’s quantum theory of radiation as a starting point, Einstein deduced that each photon must possess energy E, given by the equation :

E = hvIn which v is the frequentcy of light and h is

Planck’s constant

The equation of E = hv

E = hv E = KE + BE hv = KE + BEin which :* KE is the kinetic energy of the

ejected electron and * BE is the binding energy of the

electron in the metal

The energies that the electron in the hydrogen atom

• En = – RH

• In which RH, the Rydberg constant, has the value 2.18 x 10-18 J

• The number n is an integer called the principal quantum number; it has the value n = 1, 2, 3,….

1

n2

Strength of Covalent Bond

The Strength of Covalent Bond is defined by the amount of energy needed to break it.

A quantitative measure of stability of a molecule is its bond dissociation energy (or bond energy).

For example: H2(g) H(g) + H(g) H = 436.4 kJHCl(g) H9g) + Cl(g) H = 431.9 kJ

STRUCTURE

Covalent Bond in Organic Compounds

SATURATED: Bonding : are

formed by overlap of two atomic orbitals, each of which contains one electron

UNSATURATED : and bonding. bonding

(pi) bond

ENTHALPY

• The enthalpy of a system, which has the symbol H, is that of Heat content (heat of reaction) and is measure of the change in total bonding energy during a reaction. It is defined mathematically as : H = U + PV; H is a function of state

The standard enthalpy change for any reaction (∆H0

rxn)can determine by using standard enthalpies of formation (∆H0f) and Hess’s Law

CONSTANT PRESSURE PROCESSES

• Most processes occur in the open at one atmosphere pressure.In these cases, P1 = P2 = P, say, and

• ∆H = ∆U + P ∆V

Positive and negative sign of Enthalpy

• ∆ H has a negative sign for an exothermic change (heat is released), is mean the bonds in the products are stronger (more stable) than the bonds in the reactants.

• ∆ H has a positive sign for an endothermic change (heat is absorbed), is mean the bonds in the products are weaker (less stable) than the bonds in the reactants.

HESS’S LAW

The principle of constant heat summation, often known as Hess’s Law, is thus seen to lead directly fromthe fact that H is a function of state.

Hess’s Law be valid for : ∆r H˚ or ∆f H˚,˚ = all reactants and products are in • their standard states.f = formation standard enthalpies of formationPº = 1 atmosphere, and temperature 25ºC or 298.15°K

• This idea is immensely powerful, because it enables Hº298 values to be determined for any reaction, as long as the H of formation are known for each reactant and product.

• ∆r H = H prod – H react

Example No. 1 :

• Consider the following two chemical equations.

• 1. C(s) + ½ O2 (g) CO (g) ∆r H (1) = -110.5 kJ

2. CO (g) + ½ O2 (g) CO2 (g) ∆r H (2) = -283.0 kJ

How many Joule ∆r H (3) = ….? For below equation

C (s) + O2 (g) CO2 (g) ∆r H (3) = ...?

2 P(s) + 3 Cl2(g) 2 PCl3(l) ∆rH (1)= -640 kJ2 P(s) + 5 Cl2(g) 2 PCl5(s) ∆rH (2)= -887 kJ

Please calculate the value of ∆r H for below

equation PCl3(l) + Cl2(g) PCl5(s) ∆r H (3) = .....?• Please you make the application of Hess’s

Law, consider the use of

Example No. 2 :

2 P(s) + 3 Cl2(g) 2 PCl3(l) ∆rH (1)= -640 kJ2 P(s) + 5 Cl2(g) 2 PCl5(s) ∆rH (2)= -887 kJ

Please calculate the value of ∆r H for below

equation PCl3(l) + Cl2(g) PCl5(s) ∆r H (3) = .....?• Please you make the application of Hess’s

Law, consider the use of

solution No. 2:

A. – 247 kJ C. – 124 kJ E. – 1527 kJ. B. + 247 kJ D. + 124 kJ

Spontaneous Changes

The process tends to occur or not

Two driving forces in nature

1. The towards minimization of energy is one such directing influence, but there is also a tendency for material to become more physically disorganized.

2. The tendency for entropy to increase is nature’s second driving force.

ENTROPY

• The symbol of entropy = Sa thermodynamic function of state

NATURAL OR IRREVERSIBLE PROCESS

• The entropy of system and surroundings together increases during all natural or

irreversible process;• ∆ Ssystem + ∆ Ssurrounding = ∆ Suniverse > 0

REVERSIBLE PROCESS

• For reversible process, the total entropy is unchanged;

• ∆ Ssys + ∆ Ssur = ∆ Suniverse = 0

• For a cyclic process, a process in which the

final state is the same as the initial state,

∆S = 0

CYCLIC PROCESSES

Changes of entropy with temperature

• ∆ S = S2 – S1 = CP ln T2/T1 (P constant)

• ∆ S = S2 – S1 = CV ln T2/T1 (V constant)

Absolute entropy

The third Law of ThermodynamicsAll truly perfect crystals at absolute zero temperature have zero entropy.

FREE ENERGY

Gibbs Free Energy

The Gibbs energy determines thedirection of a SpontaneousProcess for a System at ConstantPressure and Temperature

G is function of state

Gibbs free energy, G,. It is a function of state which provides possible or not a change of any kind will tend tooccur.

The value of ∆ G

• For a favorable reaction, ∆G has a negative value, meaning that energy is released to the surroundings Exergonic

• For a unfavorable reaction, ∆G has a positive value, meaning that energy is absorbed from the surroundings Endergonic

REACTION AT CONSTANT TEMPERATURE & PRESSURE

dG ≤ 0 (constant T and P)The quantity G is called the Gibbs energy

Value of G in a system at constant T and P

• The Gibbs energy will decrease as the result of any spontaneous processes until the system reaches equilibrium, where d G = 0.

• The Gibbs free energy is defined as:

• G = H - TS

RELATIONSHIP BETWEEN THE PROCESSES WITH GIBB’S FREE ENERGY

Spontaneous processes, that is, those withnegative ∆ G values, are said to beexergonic; they can be utilized to do work.Processes that are not spontaneous, thosewith positive ∆ G values, are termedendogonic; they must be driven by the inputof free energy. Processes at equilibrium, those in which theforward and backward reactions are exactlybalance, are characterized by ∆ G = 0.

Thermodynamic

In Chemistry

System

In Living

System

What was The Thermodynamic Studied in Living System?

• Thermodynamic In Living System• 1. ∆ H (heat)• 2. ∆ S (the extent of disorder of the • system)• 3. ∆ G (Gibbs change in free energy that • proportion of the total energy change • in a system, that is available for doing • work)

Thermodynamic In Living System

• Under the conditions of biochemical reactions,

• 1. ∆ H (heat) is approximately equal

• to ∆ E, the total change in internal

• energy of the reaction,

• ∆ G = ∆ H – T ∆S, become:

• ∆ G = ∆ E – T ∆S

What is the difference between chemical reaction in nonbiologic systems and in biologic systems?

Nonbiologic systems may utilize heat

energy to perform work, but biologic

systems are essentially isothermic and

use chemical energy to power living

processes.

ATP (Adenosine Triphosphate):The Primary Energy Carrier)

• Certain bonds in ATP save the energy released during the oxidation of carbohydrates, lipids, and proteins.

• The ATP molecules act as energy carries, and deliver the energy to the parts of the cell where energy is needed to power muscle contraction, biosynthesis, and other cellular work.

Energy source in the body

Structure of ATP

ATP plays a centralrole in thetransference of freeenergy from theexergonic to theEndergonicprocesses. It serves as a carrier of chemical energybetween high energyphosphate donorsand low energyphosphate acceptors

Energy source in the body

Structure of ATP

ATP consists ofadenine (a purine), ribose and threephosphate groups,out of which thetwo terminalphosphate groupsbeing anhydridebonds are the highenergy groups

Energy change in chemical reactions

Almost all chemical reactionsabsorb or produce (release)Energy.

The standard free energy, i.e. ∆ G0’ of hydrolysis of ATP1. ATP + H20 ADP + Pi ∆ G0’ = -7.3 kcal/mol used for doing work2. ATP + H20 AMP + PPi ∆ G0’ = -7.7 kcal/mol

Chemical reaction can use up, or produce useful energy.

Exergonic reactions produce an energyoutput (ΔG = - means that the process isnot favorable)Endergonic reactions require an energyInput (ΔG = + the criterion for a favorableprocess in a nonisolated system, atconstant temperature and pressure)Biochemical system couple these energyyielding (exergonic: unstable to stable)with energy requiring (endergonic: stableto unstable) to make cellular metabolismwork.

Energy transfer in the body

• The chemical reactions within cells are accompanied by changes in energy.

• Cells accomplish their tasks by coupling energy-requiring reactions with energy-producing reactions

The working cell

Protein + ATP Pro-phosphate complex + ADP (protein is phosphorylated)

What is the reaction called, if a reaction between solute and solvent needs heat?

• A. exergonic reaction • B. endergonic reaction• C. exothermic reaction• D. endothermic reaction• E. kinetic reaction

• The reaction of glucose become to glucose-6-phosphate as follows:

• Pi + glucose glucose-6-P + H2O • ΔG0 = +13.8 (kJ.mol-1)• ATP + H2O ADP + Pi • ΔG0 = -30.5 (kJ.mol-1) ATP + glucose ADP + glucose-6-P• ΔG0 = -16.7 (kJ.mol-1)

What is the reaction above called?A. exergonic reaction B. endergonic reactionC. exothermic reaction D. endothermic reactionE. kinetic reaction

Energy transfer in the body

• Phosphoryl-transfer Reactions• R1-O-PO3

2- + R2-OH R1-OH + R2-O-PO32-

• Are of enormous metabolic significance. Some of the most important reactions of this type involve the synthesis and hydrolysis of ATP:

• ATP + H2O ADP + Pi• ATP + H2O AMP + PPi • For examples: next slide

Continuation: Energy transfer in the body

• The metabolism of glucose is its conversion to glucose-6-phosphate :

• Endergonic half-reaction 1:• Pi + glucose glucose-6-P + H2O

ΔG0 = +13.8 (kJ.mol-1)• Exergonic half reaction 2:• ATP + H2O ADP + Pi • ΔG0 = - 30.5 (kJ.mol-1)ATP + glucose ADP +glucose-6-P• ΔG0 = -16.7 (kJ.mol-1) Exergonic

Active transport: An energy-requiring process involving the movement of

substances across a membrane

• High muscle activity:1. relaxed muscle + ATP Contracted muscle +

ADP + Pi

2. ADP + phosphocreatinine ATP + creatine

• Low muscle activity:

1. Catabolic energy + ADP +Pi ATP

2. ATP + Creatine ADP + Phosphocreatine

THERMODYNAMICS OF LIVE

1. Living organism are open system and therefore can never be at equilibrium. 2. The free energy from this process is

used to do work and to produce the high degree of organization characteristic of life.

3. Living system must maintain a nonequilibrium state for several

reasons. For example: the ATP-generating consumption of glucose.

Alterations in Body Temperature

• FEVER AND HYPERTHERMIA:• Fever:• Is an elevation of body temperature

above the normal circadian range as the result of a change in the thermoregulatory center located in the anterior hypothalamus.

• A normal body temperature is ordinarily maintained, despite environmental variations, through the ability of the thermoregulatory center to balance heat oproduction by tissues (notably, muscles and the liver) with heat dissipation.

Continuation:

• With fever, the balance is shifted to increase the core temperature.

• Hyperthermia:• Is an elevation of body temperature

above the hypothalamic set point due to insufficient heat dissipation (e.g. in association with exercise perspiration-inhibiting drugs, or a hot environment) the topic in Lab activity (salicylat poisoning)

Summary References:

1. Warn, J.R.W., 1999, Concise Chemical Thermodynamics, Second Edition, Stanley Thornes Ltd., United Kingdom.

2. McQuarrie, D.A., Simon, J.D., 1997, Physical Chemistry a Molecular Approach, University Science Books, Sausalito.