Lecture notes
Thermodynamics
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- University Of Southampton (UOS)
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ThermodynamicsIntroduction of Energy
Total Internal Energy of System: U- measured in Joules
Internal Energy of a system is sum of WORK and HEAT
U= w + q w= work q= heat
Total energy of system is its internal energy (U)
Internal Energy can be transferred from one entity to another as WORK (w)/ HEAT (q)
Forms of Energy
Kinetic Energy Potential Energy
Energy a system has due to its motion2 Energy that is ‘stored’
For particle with mass (m) and velocity (v) GRAVITATIONAL ENERGY
1
Ekin= mv2 CHEMICAL BOND ENERGY- stored in chemical bonds
2
Bond formation- RELEASES energy
Bond cleavage- REQUIRES energy
Gravitational Potential Energy- energy an entity has due to its position- used to do work when an entity changes it’s
position
(Large heavy ball raised at hight position holds large amount of potential energy)
Chemical Potential Energy- energy that is stored in chemical bonds of compound- only be used when a compound
undergoes chemical change- when bonds formed/ broken
More energy is requires to separate two atoms joined by a VERY STABLE BOND- one in which free (valence) electrons
are shared in a higher- energy arrangement
Bond energy- amount of energy ‘stored’ in bond and amount of
energy required to break ionic/ covalent bond to separate two joined
atoms
Concept of System
SYSTEM- particular thing which is contained within a boundary
SURROUNDUINGS- everything else in contact with the system,
Nothing outside them changes what is within the ‘boundary’
nor is affected by the changes within the ‘boundary’
Boundary- separates system from the
surroundings
Closed system- Matter can’t be exchanged (lid on the container- prevents the contents of the container spilling
out into its surroundings)
Isolated system- (Thermos) system retains its energy- if perfectly isolated- contents of the flask stay hot or cold
indefinitely
In theory- no energy should be able to be transferred from system to surroundings- system retains its energy
,While energy is being transferred between system and surroundings- total energy of system and surroundings
ALWAYS STAY the SAME
System and surroundings to constitute a sealed environment- energy can flow between the system and surroundings
but CAN’T ESCAPE.
1st Law of Thermodynamics- CONSERVATION OF Energy
Sum of energies of system and its surroundings remains CONSTANT
Energy CAN’T be CREATED/ DESTROYED
Energy can ONLY be TRANSFERRED between different forms
Energy transfer as work (w)
WORK- Any process that can be used to lift a weight (operate against the force for gravity)
When work is performed- energy is transferred from the system to surroundings
WORK= FORCE X DISTANCE (Newton: Force= mass (m) X acceleration (a))
F=ma
Amount of energy transferred between system and the surroundings is proportional to amount of work done by
system on the surroundings
Greater the amount of energy transferred between system and surroundings- greater the amount of work
done by system on the surroundings k= Coulumbs CONSTANT
k q1 q2 q 1∧¿q ¿= magnitudes of charges, (+=
When moving charges, consider COULOMBS LAW: F= 2
2
r positive charge, -=negative charge)
r= distance between two charges
Energy associated with work- energy that can be channelled in organised
way to do something useful. E.g
To drive an ion across membrane
To power the contraction of a muscle
Energy transfer as heat (q)
Heat is the transfer of energy from hot to cold- from region of higher temperature to region of lower temperature.
Universal Observation:
Hot system in contact with colder surroundings will spontaneously get colder – energy flows as heat from system
to its surroundings
Reverse won’t happen- system with lower temperature than it’s surroundings, energy flows from surroundings
into system (HOT to COLD)
When something loses energy in this way- undergoes COOLING
Energy flows from hot object to cooler object until they both reach SAME temperature.
Transfer of energy as heat- from high temperature to lower temperature is SPONTANEOUS process
Transfer of energy doesn’t flow spontaneously from cold to hot
,Energy associated with heat- energy that is lost from hot entity to neighbouring cooler entity. E.g. warm human
body to its cooler surroundings
Heat is TRANSFER OF ENERGY from system (high temp) to surroundings (low temp). Temperature measure of
thermal energy of system- kinetic energy of molecules (motion)
Temperature of system is PROPORTIONAL to its energy
System with large amount of thermal energy- atoms and molecules have lot of kinetic energy- has HIGH
temperature
System with little thermal energy- atoms and molecules have less kinetic energy- has LOW temperature
Atoms and molecules in system don’t all possess same thermal/ kinetic energy
Some molecules possess great deal of energy
Some posses very little energy
Enthalpy
When chemical bond formed- energy is released, when chemical bond is broken- energy is consumed
If there’s more energy with formation of new bonds than with breaking bonds- reaction gives out energy
If there’s more energy with breaking of bonds than with formation of bonds- reaction consumes energy
Overall ENTHALPY CHANGE (∆H)- difference between energy consumed when bonds break and when new bonds
form.
Energy transfer during chemical reactions
Under constant pressure, enthalpy change happening in system= heat transferred to system during reaction
∆H=q
In chemical reaction, internal energy (U), may change, pressure (p) and volume (V) may change
Energy derived from pressure and volume changes: E=pV
Enthalpy (H) is sum of internal energy (U) and pv
H= U+ pV Measurements made at constant pressure:
p=pconst
Change of enthalpy of system (∆H):
Energy used when system
∆(pV)= p∆V
either contracts or expands
∆H= ∆U+ ∆(pV) during course of chemical
reaction Alternatively, with constant volume: V=Vconst
With pconstant: ∆H= ∆U+ p∆V
∆(pV)=V∆p
- If chemical reaction results in generation of smaller number of molecules that at start of reaction- system will
occupy smaller volume
Pressure being exerted on system to surroundings will squeeze smaller number of molecules- occupy smaller space.
-If reaction generates larger number of molecules that at start of the reaction- molecules occupy larger volume.
To occupy larger volume- molecules must do work on surroundings- push out against the surroundings to expand
volume that system occupies
p∆V- represents work done on system by surroundings if volume of system
, If energy required to break bonds during chemical reaction is greater than
energy released from formation of bonds- enthalpy change is positive- system
needs to supply energy for reaction to occur
energy is absorbed not system form surroundings- provide net increase in
energy required for reaction to happen
If energy required to break bond is less than energy released during formation
of bonds- enthalpy change for reaction is negative – there is excess of energy when reactants form products-
transferred from system to surroundings
Energy released by system as reaction proceeds- heat is transferred from system to surroundings
Relationship between measurement of energy change and measurable quantity (∆q): AU= ∆w+ ∆q
∆H= ∆w+ ∆q+ p∆V
Assuming, work (w) that is done by system in chemical reaction is a change in volume, V: w = -p∆V
At constant pressure: ∆H= -p∆V+ ∆q + p∆V Heat given up/ absorbed ∆q is enthalpy
= ∆q change
Endothermic and Exothermic Reactions
Exothermic reaction- energy released from system into
surroundings
Endothermic reaction- energy absorbed from surroundings
into system
If ΣEreactants is greater than ΣEproducts – more energy is required to
break bonds than is liberated from formation of bonds- ∆H is positive
If ΣEreactants is less than ΣEproducts – less energy is required to break bonds is liberated from formation of bonds-
∆H is negative
E.g. Burning of Methane
CH4 (g) + 2O2 (g) →CO2 (g) + H2 (g)
• DHreaction = Sum (Ereactants) – Sum(Eproducts)
Reactants Products
4 (C-H): 4(412)= 1648 kJmol-1 2(C=O): 2(740)= 1480 kJmol-1
2(O=O): 2(497)= 994 kJmol-1 4(O-H): 4(463)= 1852 kJmol-1
Sum (Ereactants)= 2642 kJmol-1 Sum (Eproducts)= 3332 kJmol-1
DH kJmol-1 - exothermic reaction
reaction= 2642- 3332= -690
E.g. Photosynthesis ᶱ– standard state
6CO2 (g) + 6H2O(l) →C6H12O6 (s) + 6O2 (g) 1 mole at 1 atm
• DfHᶱ = DfHᶱ - DfHᶱ
reaction (products) (reactants)
Reactants Products
o
DfH (6CO2): 6(-393.5) DfHᶱ (C6H12O6): -1271
o
DfH (6H2O): 6(-285.8) DfHᶱ (6O2): 6(0)= 0
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