Topic 4: Thermochemistry

Energy

Potential energy and kinetic energy

Heat and work

interconnected ways of transferring energy

Main Law of Thermodynamics

Energy can be transferred into different forms - IT CANNOT BE LOST

• thermodynamics describes transfer of energy from one pace to another or one form to another

Thermochemistry units

Photosynthesis and Cellular resp

The first law of thermodynamics

Three types of system

open = matter and heat transfer bw system and surrounding

closed = no matter but heat transfer occurs

isolated = heat transfer and matter transfer prevented

Isothermal exchange

heat transfer from hot body to cold body until temps are the same

Adiabatic change

system and surrounding have no heat exchange so temp might not be equal

Heat (q) and temp not same thing

heat: quantifiable energy

temp: measure of how hot smth it (kinetic energy)

• eg. both larger and smaller coffee has same temp but larger coffee has more heat

extensive property

property of matter that changes depending on amount of matter

• eg. heat

intensive property

prop. of matter that does not change if amount of matter changes

• eg. temp

Work (w)

energy exchange due to motion against an opposing force

• when system does work on surroundings - it LOSES energy

• when surrounding do work on system - it GAINS energy

Work related gases - eq. for change in volume

State functions

we only know final end result not how we got to end result

• eg. internal energy (U)

Internal energy (U)

sum of all its kinetic and potential energies

• total energy

Delta U in closed system

if positive - we gaines energy

if negative - we lost energy

if 0 = no change

How can we change Delta U

Electrical motor example for change internal energy

Issue with internal energy

some of the enrgy transferred to the system is used to do the work of expansion against constant pressure

• in order to not have to always worry abt this - enthalpy (H) is used

Enthalpy (H) and enthalpy change

Endothermic vs exothermic

Energy profile diagrams

Delta H and Delta U relationships

Quick example on delta U and Delta H calc.

Melting enthalpy changes

absorb energy bc for bonds to break - the potential energy has to increase

freezing is opposite

Vaporisation enthalpy changes

endothermic

How to approach this

Heat capacities

how much heat energy required to raise the temp by 1deg

Question from before now using heat capacities

Heat capacities in gases

Cp is generally Cv + R

energy profile mechanism

when reaction happens - enthalpy either decreases or increases

Why kind of property is enthalpy change

extensive property - depends on amount of substance

Enthalpy change - thermochemical reactions

Calculating enthalpy change

H(products) - H(reactants) or

bonds broken - bonds form

Solution calorimetry

Bomb calorimetry

bc v doesn’t change - heat is under constant voluem/pressure

Bond calorimeter calculations

heat capacity is specific to calorimeter (eg. CF) and bc constant volume → Delta U = Delta H

What can enthalpy be impacted by

pressure, conc of reactants and products, phyisical state and form

• we need to define standard states

Standard enthalpy change of reaction

What standard states are in example

Hess’s law

basically if we know the enthalpy changes of half the reactions like step by step - we can just add them together to get the total enthalpy change bc it’s a state function

Hess’s law written

Things we have to consider when applying Hess’s law

must manipulate

Example of Hess’s law applied

reaction 1 is fine and we can use that delta H value

Reaction 2 adjustment

Reaction 3 adjustment

Final check to do before adding the adjusted eqs tgt

everything must cancel out to give you og reaction

Standard enthalpy change of formation

when elements in their standard state - eg. H2(g)

using tabulated values to find actual delta H

Formula to find the delta r H

(coeff x product’s H value from table) - (coeff x reactant’s H value from table)

example of standard enthalpy change of formation

O is 0 bc its in standard state

Bond enthalpies

energy required to break a bond is the bond enthalpy

• the process of breaking bond = homolytic dissociation

The actual equation to find delta H via bond enthalpies

Different ways to calculate delta H

System’s enthalpy depends on temp

when temp increases, the enthalpy of substance increases

The Kirchhoff equation

always products - reactants when deltas

Example of Kirschhof eq.

Spontaneous processes

process, which once started, occur without any external intervention or action (may need energy to start tho)

• spontaneous direction of reaction goes toward equilibrium

• usually exothermic

Are spontaneous reactions always exothermic?

no - eg. dissolution

• they can be endothermic

Do all spontaneous reactions have delta H

no - eg. gases mixing - no heat exhcange

Conditions that spontaneous processes depend on

delta H but also temp

• not related to ROR

• ENTROPY IS IMPORTANT

Entropy (S)

measure for the degree of disorder in a system, its surrounding and universe

• eg. shuffling increases entropy - likelihood of getting back to og is very low

Entropy in terms of second law of thermodynamics

if Delta S is + → more disorder

if Delta S is - → less disorder

molecular distribution

middle two more likely

microstates and macrostates + higher entropy

most probable is the macrostate with the highest number of microstates

• eg. middle

middle one has higher entropy

Example of macrostates with 4 particles prob

What happens as the no. of atoms in gases increase

• more an more microstates

• less likely to find all on left

• chances of surprising observations decrease

• when number of atoms large, we are at the thermodynamic limit

Thermodynamic limit

Boltzmann equation allows us to calculate entropy

Example of using Boltzmann equation

Understanding entropy

we want to maximise a systems entropy

• this is the most probably distribution of atoms wher matter and energy is MOST DISPERSED

• this is why gases are so disordered and dispersed

Arrow of time

Reversible processes

system and surroundings can be put back to g states

• eg. entropy doesn’t change in reversible system

• temp can leave system and come back in

• state changes

• maximum amount of work is obtained

Third Law of Thermodynamics

entropy increases with temp

• can get a value of the entropy of water bottle and by what bc we have a reference point

How can we find entropy

How to calc entropies

the heat put in in joules divided by temp

Why hotter water will have lower entropy value than cold

the hotter water was already more disordered so adding heat changes entropy less than cold water does

• therefore, colder water have higher entropy increase

Why does heat always flow from hotter part to cooler part?

if system loses heat (so left bc the temp is already high) = minus sign

if it gains heat = plus sign

We can find the molar entropies of substances at certain temps

lower temp = lower entropy

Standard entropy change of reaction

Entropy on molecular scale

1. Bigger/more complex molecule = higher entropy More atoms = more ways to move, rotate, vibrate = more disorder

2. Heavier molecule = higher entropy More mass = more accessible microstates

Entropy by states

gas > liquid > solid bc more freedom of motion

• eg. H2 + O2 → H2O (l) = entropy decreases

Phase changes and Temp increase

at the exact state change points (phase changes = melting point/ evaporation point) → very sharp

• entrop increases bc temp is being added to system

How to calculate entropy at points of phase change

Phase changes and equilibrium

At a phase change (boiling or melting), you can calculate ΔS if you know ΔH and the temperature.

• For boiling: ΔS = ΔH_vap / T_boiling

• For melting: ΔS = ΔH_fus / T_melting

Enthalpy in J and temp in K

eg."Given enthalpy of vaporisation of mercury is 58.51 kJ/mol and boiling point is 629.7K, calculate ΔS"

ΔS = 58510 J/mol ÷ 629.7K = 92.9 J/K/mol

Law 2 thermodynamics

Entropy calc if given heat capacity

Determining whether a reaction is spontaneous - calculations incl. delta H and S

Q asks: delta S universe

Step 1: delta S system by

entropies of products - entropies of reactants

Step 2: delta H system by

Enthalpy changes of formation → products - reactants

Step 3: delta s surroundings by

-delta H system / T in K

Step 4: delta S universe

add both values tgt and i positive then spontaneous

How to read the final Delta S universe value in terms of spontaneity

special cases: is exothermic and negative entropy → compare magnitude

vice versa

Gibbs Free Energy Change

What exactly is Gibbs free energy (G)

bc S and H are state functions, so it G

bc T is always positive, if S is positive then multiply → has to be negative

Example of determining spontaneity of reaction

Delta G with polymerisation

at low temps, delta H is negative but at high temps -TAS will dominate and delta G becomes positive

• polymerisation no longer proceeds at certain temps

Why ice melts and water freezes spontaneously

Phase changes worked example

Crossovers between enthalpy, entropy and gibbs

Standard gibbs free energy changes (Delta f G)