Doc Brown's A Level Chemistry Advanced Level Theoretical Physical Chemistry – AS A2 Level Revision Notes – Basic Thermodynamics
GCE Thermodynamics–thermochemistry sub–index links below
Part 3: ΔS Entropy Changes and ΔG Free Energy Changes
3.2 Table of entropy data & comments and 3.3 Entropy values, entropy changes and feasibility of a physical change
Having introduced the idea of entropy this section looks at the entropy values (S) of variety of elements and compounds AND in some cases their entropy in different physical states and then discusses why they are different i.e. why is one molecule gat a higher/lower entropy than another and then the entropy change for physical process like melting are described. The total entropy change for the 'system' and its' surroundings' must be considered.
Energetics index: GCSE Notes on the basics of chemical energy changes – important to study and know before tackling any of the three Advanced Level Chemistry pages Parts 1–3 here * Part 1a–b ΔH Enthalpy Changes 1.1 Advanced Introduction to enthalpy changes – reaction, formation, combustion : 1.2a & 1.2b(i)–(iii) Thermochemistry – Hess's Law and Enthalpy Calculations – reaction, combustion, formation etc. : 1.2b(iv) Bond Enthalpy Calculations : 1.3a–b Experimental methods for determining enthalpy changes and treatment of results : 1.4 Some enthalpy data patterns : 1.4a The combustion of linear alkanes and linear aliphatic alcohols : 1.4b Some patterns in Bond Enthalpies and Bond Length : 1.4c Enthalpies of Neutralisation : 1.4d Enthalpies of Hydrogenation of unsaturated hydrocarbons and evidence of aromatic ring structure in benzene : Extra Q page A set of practice enthalpy calculations with worked out answers ** Part 2 ΔH Enthalpies of ion hydration, solution, atomisation, lattice energy, electron affinity and the Born–Haber cycle : 2.1a–c What happens when a salt dissolves in water and why? : 2.1d–e Enthalpy cycles involving a salt dissolving : 2.2a–c The Born–Haber Cycle *** Part 3 ΔS Entropy and ΔG Free Energy Changes : 3.1a–g Introduction to Entropy : 3.2 Examples of entropy values and comments * 3.3a ΔS, Entropy and change of state : 3.3b ΔS, Entropy changes and the feasibility of a chemical change : 3.4a–d More on ΔG, Free energy changes, feasibility and applications : 3.5 Calculating Equilibrium Constants : 3.6 Kinetic stability versus thermodynamic feasibility * PLEASE note that delta H/S/G values vary slightly from source to source, so I apologise in advance for any inconsistencies that may arise as I've researched and developed each section.
3.2 Table of entropy data and comments
The examples below illustrates some of the ideas described in section 3.1 above.
S denotes the entropy content of a substance and it always has a positive value. It increases with increase in temperature as more energy levels (rotational, vibrational, electronic etc.) become more available as energy is absorbed and redistributed as the temperature increases.
Extra comments and ideas
3.3 ΔS, Entropy changes and the feasibility of a change – physical or chemical
3.3a Entropy Changes of State Changes
The direction of change of a system which is not at equilibrium (physical or chemical) is determined by the overall entropy change. Upto now we haven't considered the whole situation where there is potential for change. The whole situation meaning system and surroundings.
From data tables of entropy (e.g. as in table 3.2 above) you can readily calculate the entropy change of the system.
The entropy change for the surroundings is minus the enthalpy change of the system change divided by the absolute temperature in Kelvin.
The total entropy change is then given by the 'simple' equation i.e. the sum of all the entropy changes.
For a change to be feasible the total entropy change must be at least greater or equal to zero.
In this section 3.3a we will apply this rule to two physical state changes.
3.3a1 The melting of a pure substance – water
3.3a2 The freezing of a mixture e.g. salt solution
The spreading of salt (sodium chloride) on roads is a tried and tested method of de–icing roads or preventing ice formation. So, why does it work? why is the melting/freezing point of water lowered? – well its all down to the entropy change! Unfortunately I can't find accurate enough data to do an illustrative calculation. However, using the entropy equations, we can do a 'thought experiment' and reason out why adding salt lowers the freezing point of water, or to state it another way, why salt solution has a lower freezing point than pure water.
The most important point to grasp is that salt solution has a greater entropy than pure water. There are now three different particles in the liquid instead of one i.e. sodium and chloride ions as well as water molecules. When salt solution freezes pure water will crystallise out of the liquid.
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