ENTROPY. Definition: Entropy is the quantitative measure of disorder in a system

Transcription

1 ENTROPY Definition: Entropy is the quantitative measure of disorder in a system This depends not only on the degree to which the molecules are randomly arranged, but also on the random distribution of energy quanta between molecules Essentially, for a process to occur spontaneously the degree of disorder/ dispersion must increase This is the 2nd Law of Thermodynamics

2 Lego = High Entropy = Nightmare Low entropy toy In a closed system, natural spontaneous change involves the increase in TOTAL entropy But what do we mean by spontaneous here.? A spontaneous process takes place without any continuous intervention from outside the system NB. The process may need a kick start to get going! 2

3 Distribution and disorder If we let the sheep roam around and then take an aerial picture of the field, the black and white sheep would be randomly distributed: The natural tendency of the universe is to disorder! 3

4 * Suppose the farmer wants to separate the sheep. A sheepdog will have to do work to separate them: * Not only that, but the dog will have to keep working to keep them separated, because their natural tendency would be to randomly distribute throughout the field * We can postulate that if energy must be expended to keep the sheep from being randomly distributed, then if we reverse the process (i.e. start with sheep in an ordered arrangement and let them become randomly distributed) then energy is somehow released 4

5 Energy and disorder Energy comes in packets called quanta. The more quanta, the greater the number of possible distributions of the quanta and therefore the greater the entropy. section=tutorial&article=8 Diffusion Diffusion is a good example of a spontaneous process - the system will adjust in order to maximise entropy, but be careful! It will maximise the TOTAL ENTROPY which is not just dependent on the disorder of the system (see later). 5

6 Entropy of substances (under given conditions) What do you notice about the Standard Molar Entropies? The entropy of liquid water is greater than the entropy of solid water (ice) at 0 C. 6

7 Zero ENTROPY The entropy of a perfect crystal at absolute is zero This is the 3rd Law of thermodynamics Good website: chemed.chem.purdue.edu/genchem/ topicreview/bp/ch21/entropy.php What is the 1st law of Thermodynamics? What is the Zeroth law of Thermodynamics? 7

8 The Zeroth law tells us that temperature is the key to knowing how heat will transfer between 2 systems It doesn t matter how much energy they have in total It established temperature as the fundamental and measurable property of matter This is why it was given the top spot in the laws. It precedes the others. Systems that are in thermodynamic equilibrium are at the same temperature

9

10 Try to get into your head an idea of which substances have low entropy and which have high. Remember some examples!! In general: Gases have high entropies (why?) - reactions that produce gases probably lead to an increase in entropy Ionic substances increase their entropy significantly when they dissolve (why?) Breaking down large molecules increases entropy (why?) 10

11 Change in entropy of the system - ΔSsys When considering a process, ask yourself: Does the number of possible arrangements of the particles present increase (and does the number of arrangements of the energy quanta increase)? Remember: A reaction is considered to be spontaneous if it occurs without external influence (irrespective of rate). Up to now in the course we have only considered Thermodynamic and Kinetic factors. Now we must appreciate the importance of ENTROPY 11

12 Use ideas of entropy to explain what is happening when Sodium Chloride dissolves in water

13 Talk about ENTROPY in reaction 1 in relation to the number of moles in the reactants and products and the states of each

14 Explain the shape of the following graph in terms of Entropy Textbook pp62-63

15 Now let s get quantitative To determine whether a change is spontaneous, we must not only consider the change in entropy of the substances involved (the system), but also in the surroundings ΔStotal = ΔSsys+ΔSsurr This is why endothemic reactions can be spontaneous, and exothemic don t have to be! ΔSsys is given by 15

16 ΔStotal = ΔSsys + ΔSsurr ΔSsurr This is calculated by the equation: ΔSsurr = -ΔH/T Where ΔH is the enthalpy change for the reaction (in J mol -1 ) T is the absolute temperature (K) So, What can we say about ΔSsurr for exothermic and endothermic processes? 16 Look at the worked examples on page 65

17 17

18 ΔStotal = ΔSsys+ΔSsurr Dissolving ammonium nitrate in water ΔStotal = ΔSsys - ΔH/T Equation: This is an spontaneous endothermic change at 298K - try and explain why?

19 Consider the following reaction Write the state symbols below each of the species involved This is an endothermic change which is spontaneous at room temperature because the increase in entropy due to the gaseous carbon dioxide formed outweighs the fall in entropy of the surroundings.

20 The reaction between magnesium ribbon and oxygen What is the total entropy change if a reaction at 293 K? The enthalpy change for the reaction is kj mol -1, and the entropy change of the system is -216 J K -1 mol -1. There has been a very large increase in entropy but the reaction is not visibly spontaneous at room temperature. We say that the reaction is feasible but you need to supply a lot of activation energy to get it to burn. We cannot assume that a feasible reaction with ΔStotal will happen at a particular temperature. There are other factors involved. Textbook pp64-69

21 For a process to be thermodynamically feasible ΔG must be negative. If ΔG is positive the process is not feasible If ΔG is ZERO the reaction is in equilibrium. Try to fill in the following table of possible scenarios ΔS Positive Negative ΔH Positive ΔG negative if magnitude of TΔS is larger than that of ΔH ΔG always positive 21 Negative ΔG always negative ΔG negative if magnitude of TΔS is smaller than that of ΔH

22 Feasibility and Spontaneity ΔG = ΔH - TΔSsys Remember, For a process to be thermodynamically feasible ΔG must be negative. If ΔG is positive the process is not feasible If ΔG is ZERO the reaction is in equilibrium. However, this does not mean that the reaction will proceed at an appreciable rate. There are 3 caveats 1. Kinetics: If the activation energy for the reaction is high, the rate of reaction may be very low 2. The reaction may need a kick start of energy, and then we can use the phrase the reaction will proceed without any continuous influence from outside the system 3. Non-standard conditions ΔGº is calculated for 298K and 100kPa, with and 22 solutions of 1M concentration. Changing concentration and temperature can alter the favoured direction of a reaction. The reaction of Magnesium with Oxygen or Hydrogen or Oxygen excellently illustrates points 1 and 2.

23 Worked example - Decomposition of Calcium Carbonate (p71) Copy the standard molar entropies for the substances involved below CaCO3 CaO CO2 Hence, calculate ΔS Hence, calculate ΔG at 25ºC

24 What is the minimum temperature at which Calcium Carbonate ceases to be thermodynamically stable, IE. its decomposition becomes spontaneous? Textbook pp70-71

25 Feasibility and Equilibrium Constant ΔG º = -RTlnK R is the gas constant with a value of J K -1 mol -1. T is the temperature of the reaction in Kelvin. K is the thermodynamic equilibrium constant and has no units ΔG This is the standard free energy change here - NOT the free energy change at whatever temperature the reaction was carried out (remember that the position of equilibrium and therefore K are affected by temperature). NB - ΔG is given in kj per mole and you must convert it to J per mole before you use it in the above equation. Challenge: Rearrange the equation to make K the subject

26 What is the relationship between ΔG and K? Consider our old favourite, the Haber process If ΔG for a reaction very large, then this implies that the reaction either almost doesn t happen or that it almost goes to completion. This is nicely illustrated by the graph on p73

27 K and temperature ΔG º = -RTlnK and ΔG = ΔH - TΔSsys So, lnk = Discuss the effects of the sign of ΔH on K Textbook pp72-73

28 Extra curricular - The Solubility of salts in terms of Entropy ΔG = ΔH - TΔSsys Look at the worked examples on pp74-75 regarding solubility. What can we deduce as general rules for solubility in terms of ΔH and ΔSsys and their relative affects on ΔG?

29 Strong and weak acids We have said before that HF is a much weaker acid than Hal due to the relative strengths of the bonds, but this is not the full story Read p76-77 and try and answer the following Why is HF a weak acid and HCl is considered to be fully dissociated? Why does Ka increase as the number of Cl atoms attached to ethnic acid increase? Exam Style Questions

30

31