Introduction to

thermochemistry

Heat, work, energy and the First

Law

Learning objectives

• Define energy and identify types of energy

• Compare and contrast heat and work

• Describe internal energy and how it changes

during a process

• Describe basic properties of state functions

• Apply first law of thermodynamics to

determine heat flow and work

• Define enthalpy

Behind it all

• Why do chemical changes

happen?

• Substances spontaneously

move towards greater

stability – in energy terms

(see later how this is defined)

• High energy state is unstable

with respect to lower energy

state

• Simple (but incomplete)

analogy is ball rolling

downhill

Energy

• Is capacity to do work

• Mechanical work is application of force over distance

• Heat is energy transferred by virtue of temperature gradient – associated with molecular motion

• Joule demonstrated experimentally that heat and work are interchangeable forms of energy

F x d

Energy: forms

• Kinetic energy is energy of motion

• Potential energy is energy stored – by

position, in spring, in chemical bond, in

nucleus

2

2

1mvEK

mghEP

Energy: units

• From definition of kinetic energy (1/2mv2), we

get units of energy:

kg m2/s2

• S.I. unit for energy is joule (J) = 1 Nm

• Another common unit is calorie (cal):

Energy required to raise temperature of 1 g

water 1ºC

1 cal = 4.184 J

• Note the food calorie (Cal) = 1 000 cal

Interchange and conservation

• Energy can be changed from one form to another – Stationary ball on hill

has potential energy (P.E.) by virtue of position but no kinetic energy (K.E.). As it rolls down, it gains K.E. at the expense of P.E.

Energy conservation • There is no gain or loss:

Energy cannot be created or

destroyed; it can only be

changed from one form to

another

– Chemical processes involve

conversion of chemical

potential energy into other

forms and vice versa

– Energy never goes away, but in

some forms it is more useful

than others

– Efficient energy use means

maximizing the useful part and

minimizing the useless part

Some like it hot

• Thermal energy is the kinetic energy of

molecular motion

– Temperature measures the magnitude of the

thermal energy (all molecules at same T have

same E)

• Heat is the transfer of thermal energy from a

hotter to a cooler body

– Temperature gradient provides the “pressure” for

heat to flow

• Chemical energy is the potential energy

stored in chemical bonds

System and surroundings

• Any process can be divided into SYSTEM

contained within SURROUNDINGS

– When energy changes are measured in chemical

reaction:

– system is reaction mixture

– surroundings are flask + room + rest of universe

Internal energy (U or E)

• Internal energy is sum of all types of energy

(kinetic and potential) of system. It

measures capacity to do work

• Typically we don’t know absolute value of U

for system

– Internal energy usually has symbol U. Other

sources use E (I know, that is confusing)

• We measure change to internal energy

initialfinal UUU

Work and internal energy

• Work done on system increases its

internal energy

• Work done by system decreases its

internal energy

• ΔU and w have same sign

ΔU = w

Workin’ for a livin’

• Mechanical work is

force applied over a

distance

W = F x d

• In chemical process

release of gas

allows work to be

done (PΔV)

Work done at constant pressure

• Gas generated in reaction pushes against the piston with force: P x A

• At constant P, volume increases by ΔV and work done by system is:

w = -PΔV (ΔV = A x d) – Work done by system is –ve in expansion (ΔV > 0)

• ΔU < 0 (ΔV > 0, -PΔV < 0)

– Work done by system is +ve in contraction (ΔV < 0)

• ΔU > 0 (ΔV < 0, -PΔV > 0)

VPw

Expansion work

• Work done by gas expanding:

w = -PexΔV • In expansion the ΔV > 0; w < 0

Internal energy decreases ΔU < 0

• In contraction, ΔV < 0; w > 0

Internal energy increases ΔU > 0

Deposits and withdrawals

• Process is always viewed from perspective of system

• Energy leaving system has negative sign – (decreases internal energy – lowers energy bank balance)

• Energy entering system has positive sign – (increases internal energy – increases energy bank balance)

• Change in internal energy due to heat and/or work

wqU

First Law of Thermodynamics

Total internal energy of isolated system is constant

– Energy change is difference between final and initial states (ΔU = Ufinal – Uinitial)

– Energy that flows from system to surroundings has negative sign (Ufinal < Uinitial,)

– Energy that flows into system from surroundings has positive sign (Ufinal > Uinitial.)

Significance of state functions

• State Function A property that depends only on present

state of the system and is independent

of pathway to that state

• Internal energy is a state function, as

are pressure, volume and temperature

• Any function made from other state

functions is also a state function

• Change in state function between two

states is independent of pathway

• Given two states of system:

– ΔU is always the same

– q and w depend on type of change

Heat and work

• Any chemical process may have heat

and work terms

• Total internal energy change = sum of

contributions from each

• In closed system ΔV = 0, so q = ΔU

VPqwqU

VPUq

Cracked pots and enthalpy

• Most reactions are conducted in open

vessels where P is constant and ΔV ≠ 0

• The heat change at constant pressure is

• Enthalpy (H) is defined as:

VPUqP

PVUH

Heats of reaction and enthalpy

• Absolute enthalpy of system is not known

• Enthalpy change is measured

• Enthalpy change is known as heat of reaction

– If reaction is exothermic and involves

expansion: • ΔU < 0, ΔV > 0 ΔH less negative than ΔU

• Enthalpy change is portion of internal energy available as heat after work is done by system to expand

• If no work done, all internal energy change is enthalpy

VPUqH P

Comparing ΔH and ΔU

• In reactions involving volume

change at constant finite P, ΔH and

ΔU are different. How big is it?

• Consider reaction:

• 1 additional mole of gas is produced

• Work done by system, w < 0

ΔU = - 2045 kJ, ΔH = - 2043 kJ

PΔV = + 2kJ (P = 1 atm, ΔV = 20 L)

)(4)(3)(5)( 22283 gOHgCOgOgHC

A word on units: PV work and

joules