Halesworth U3A Science Study Group

ThermodynamicsOr

Why Things are How They Are

Or

Why You Can’t Get Something For Nothing

Ken DerhamIncludes quotations from publicly available internet sources

Halesworth U3A Science Group

What is Thermodynamics

� Thermodynamics is the branch of physics concerned with heat and temperature and their relation to energy and work.

� It defines macroscopic variables, such as internal energy, entropy, and pressure, that describe a body of matter or radiation.

� Thermodynamics applies to a wide range of topics in science and engineering, including:

� Physical chemistry

� Chemical engineering

� Mechanical engineering

Thermodynamics in a nutshell

• Energy is conserved

• It tends to disperse

• Things become more random

History 1650s-1670s

Halesworth U3A Science Study Group

Hooke & Boyle – Boyle’s Law

PV/T=c

or

P1V1 = P2V2

T1 T2

History - 18th & early 19th c

� 1750s- Joseph Black - Distinction between heat

and temperature

� 1820s – Sadi Carnot - worked on efficiency of

heat engines

History – later 19th c

� 1850s – William Thompson (Lord Kelvin) –

coined the term thermo-dynamics, linking heat

and power, and was first to formulate concise

definitions of thermodynamics

� 1870s – J Willard Gibbs – developed analysis of

energy, entropy, volume, chemical energy, etc to

determine whether a process would occur

spontaneously

What is Heat?� Until the mid-19th century heat was thought of as a

“caloric” fluid.

� The Scot, James Clerk Maxwell & the Austrian Ludwig

Boltzmann understood that a hot substance one in

which its atoms move quickly.

� The heat of an object is the total energy of all the

molecular motion inside that object.

� Temperature, on the other hand, is a measure of the

average heat or thermal energy of the molecules in a

substance

Equilibrium

� Many important practical engineering applications,

such as heat engines, refrigerators etc can be

approximated as systems consisting of several

subsystems at different temperatures and pressures

but in equilibrium.

• If two systems are each in thermal equilibrium with a

third, then they are also in thermal equilibrium with

each other

First Law of Thermodynamics

� The increase in internal energy of a closed system is

equal to the difference of the heat supplied to the

system and the work done by the system:

� ΔU = Q –W

� The internal energy of an isolated system obeys the

principle of conservation of energy; i.e.

� Energy can be transformed (changed from one form to

another) but cannot be created or destroyed

Enthalpy

� Enthalpy is a measurement of energy

� It includes the internal energy (U), which is the energy required to create a

system, and the amount of energy required to make room for it by

displacing its environment and establishing its volume and pressure

� The enthalpy of a homogeneous system is defined as:

H = U + pVwhere

H is the enthalpy of the system,

U is the internal energy of the system,

p is the pressure of the system,

V is the volume of the system

Second Law of Thermodynamics

�Heat flows from a hotter location to a

colder location.

� There are many versions of the second law, but they

all have the same effect, which is to explain the

phenomenon of irreversibility of nature.

Second Law of Thermodynamics

� Kelvin’s research into the nature of heat led to his

formulation of the second law of thermodynamics,

which states that that

Heat will not flow from a colder body to a

hotter body.

� It was first formulated to explain how a steam engine

works.

Second Law of Thermodynamics

� Kelvin's statement of the law says that heat from a

high-temperature energy source cannot be entirely

converted to 'work'. Some of the heat will be

reduced to low-quality energy and 'lost' to the

process.

� This proved that it is impossible to have a heat engine

that is 100% efficient.

Entropy

A measure of the energy that is not available for work during a thermodynamic process.

Entropy

A measure of the randomness of the microscopic constituents of a thermodynamic system.

Symbol: S

The Second Law & Entropy

� Entropy is a measure of the level of disorder of a system.

� Although it's difficult to measure the total entropy (S) of a system, it is fairly easy to measure changes in entropy (ΔS).

� For a thermodynamic system involved in a heat transfer of size Q at a temperature T, a change in entropy can be measured by:

ΔS = Q / T

The Second Law & EntropyThe second law of thermodynamics can be stated in terms of entropy.

If a reversible process occurs, there is no net change in entropy.

In an irreversible process, entropy always increases, so the change in entropy is positive.

The total entropy of the universe is continually increasing.

The Second Law & Entropy

Mathematically:

ΔS ≥ 0

I.e. the change in entropy is always greater than or equal to zero

Entropy

� (in cosmology) a tendency for the universe to attain a state of maximum homogeneity in which all matter is at a uniform temperature (heat death)

Time� In any process in which heat exchange does not occur (or

when the heat exchanged is negligible) we see that the

future behaves exactly like the past.

� E.g. In the motion of the planets in the solar system heat

is almost irrelevant. The same motion could equally take

place in reverse without any law of physics being

infringed.

� As soon as there is any transfer of heat, the future is

different from the past.

Time� As soon as there is any transfer of heat, the future is

different from the past.

� E.g. If there were no friction a pendulum can swing

forever. If we filmed it and ran the film in reverse we

would see no difference.

� But there is friction, so the pendulum heats its supports

and surroundings slightly, loses energy and slows down.

Immediately we are able to distinguish the future

(towards which the pendulum slows) from the past.

Heat and Time

� The difference between the past and the future only

exists when there is heat.

� The fundamental phenomenon that distinguishes the

future from the past is that heat passes from things

that are hotter to things that are colder.

Why does heat move from hot things to cold things

and not the other way?

� It is sheer chance!

� Boltzmann showed that it is statistically more

probable that a quickly moving atom of a hot

substance collides with a cold one and passes on a

little of its energy. Energy is conserved in collisions,

but tends to get distributed in more or less equal

parts when there are many collisions.

Why does heat move from hot things to cold things

and not the other way?

� In this way the temperature of objects in contact with

each other tends to equalise.

� It is not actually impossible for a hot body to become

hotter through contact with a cooler one, it is just

extremely improbable.

The Kelvin temperature scale

� Kelvin realised that it would be useful to be able to

define extremely low temperatures precisely.

� He noted that molecules stop moving at absolute

zero. In 1848, he proposed an absolute temperature

scale – now called the 'Kelvin scale' – where absolute

zero is 0 kelvin (0 K).

� Absolute zero on the Kelvin scale = minus 273.15

degrees on the Celsius scale.

� On the Celsius scale, water freezes at 0 degrees. On

the Kelvin scale, it freezes at 273.15 kelvin.

Third Law of Thermodynamics

� As a system approaches absolute zero, the entropy

of the system approaches a minimum value

or

� The entropy of all systems and of all states of a

system is the smallest at absolute zero

Or equivalently:

� It is impossible to reach the absolute zero of

temperature by any finite number of processes

Gibbs Free EnergyWillard Gibbs, 1873, defined a thermodynamic quantity equal to the enthalpy (H) of a system or process, minus the product of the entropy (S) and the absolute temperature (T)

G = H – TS

G is known as the Gibbs Free Energy or Gibbs Energy

Gibbs Free EnergyIn chemical reactions the change in free energy (at

constant temperature) is expressed as

ΔG = ΔH − T ΔSchange in free energy change in enthalpy (temperature x) change in entropy

If ΔG<0 reaction will be spontaneous

If ΔG=0 the system is at equilibrium (reversible)

If ΔG>0 the process will not be spontaneous and would require an

input of energy to occur

How does thermodynamics help us in

our daily lives?

� “Because refrigerators, car engines and power plants are thermodynamic machines."

� That’s true, but you don't have to understand thermodynamics in order to know those things. You can simply accept them and that's that.

� Thermodynamics is so much more than that.

How does thermodynamics help us in our daily lives?

� Thermodynamics provides a framework in which the

universe operates. In other words, anything you are

likely to encounter in your daily experience can be

broken down to thermodynamics.

� Anything.

� Therefore, understanding the fundamental laws of

thermodynamics is a fundamental part of being a

rational being.

How does thermodynamics help us in our daily lives?

� Conservation of energy is everywhere.

� This law can (almost) never be broken.

� Anything that happens around you, happens for a

reason, and that energy is not simply created.

� The way the universe operates is easier to grasp if

you search for answers under the first law of

thermodynamics.

How does thermodynamics help us in our daily lives?

� In fact, you can't even break even.

� If you have a limited amount of energy, its availability

will decrease with time.

� This is the second law of thermodynamics.

� This means that natural systems are inherently

inefficient, and there is no easy way to overcome this

problem (we can't just simply make a 100% efficient

engine).

How does thermodynamics help us in our daily lives?

� If people reason based on the laws of

thermodynamics, misinformation would not be as

widespread as it is today.

� I.e. misinformation from economics, to politics, to

engineering, to science in general.

� People tend to believe in the most absurd things,

some of which are easily proven wrong if people

knew the laws of thermodynamics.

Application to Pharmacy & Pharmacology

Every aspect of how a drug behaves in solution and (more

importantly) within the body is governed by the simple (?)

principles of thermodynamics. Drug solubility, partitioning

between immiscible solvents and drug – receptor binding can all

be understood based upon the description of such systems

according to thermodynamic terms. In fact, our understanding of

these properties is critically dependent upon a basic

understanding of the three fundamental laws of

thermodynamics.

8 minute video by a professional

Professor Dave

https://www.youtube.com/watch?v=8N1BxHgsoOw

All you need to remember:

• Energy is conserved

(Energy cannot be created or destroyed)

• Energy becomes more dispersed

(High temperature/high concentration of energy

tends towards lower temperature/dispersed energy)

• Things become more random

(Increase in disorder)