Chapter 27. Energy and Disorder

Transcription

1 Chapter 27 Energy and Disorder

2 Why Reactions Occur Exothermic Rxns - Take place spontaneously Go from high energy to low energy Downhill Endothermic Rxns. - Not usually spontaneous Go from low energy to high energy Uphill

3 Why Reactions Occur Natural Processes tend to: 1. Go from a state of high energy to a state of lower energy 2. Go from an orderly state to a disorderly state There are some exceptions to this general rule.

4 Why Reactions Occur Unless otherwise stated, rxns. will take constant temp. & constant pressure. Isothermal Processes - rxns. taking constant temp. Isobaric Processes - rxns. taking constant press. Themodynamics - studies concerning the flow of energy

5 State Functions Change in temp. - ΔT = T 2 - T 1 T 2 - final temp. T 1 - initial temp. ΔV = V 2 - V 1 ΔP = P 2 - P 1

6 State Functions State Function - one whose value depends only on the current state of the system Ex) T, P, V Amt. of change in a state function depends only on initial and final states Does not depend on path from initial to final state.

7 Internal Energy - U State function - only interested in ΔU Every system has some internal energy 2 ways of transferring energy to a system: 1. Heating the system 2. Doing work on it

8 Internal Energy - U A syst. may also transfer energy to its surroundings by giving off heat or doing work on surroundings Change in energy of a syst.: ΔU = q + w q = amt. of heat absorbed by syst. w = amt. of work done on syst.

9 Internal Energy - U q & w are not state functions depend on path followed in getting from 1 state to another q = (+) if heat flows into syst. q = (-) if heat flows out of syst. w = (+) when surrounding do work on syst. w = (-) when syst. does work on surroundings

10 Internal Energy - U When a piston compresses gas in cylinder, it does work on syst. Several ways of doing work most important to chemists is expansion or compression of syst. - press - vol work If syst. expands, it does work on its surroundings energy is being transferred from syst.

11 Enthalpy Rearrange eqn. - q = ΔU - w Since most work in lab is press-vol const. press, w = -PΔU q p = ΔU - (-PΔV) or q p = ΔU + PΔV q p = Δ(U + PV) p - const. press. process

12 Enthalpy Enthalpy - The quantity U + PV given symbol H q p = ΔH Enthalpy - heat content - is a state function ΔH = H 2 - H 1

13 Enthalpy Exothermic rxn. - products have less enthalpy than react. ΔH < 0 (is -) H f < H i Endothermic rxn. - products have more energy than react. ΔH >0 (is +) H f > H i

14 Enthalpy Change in enthalpy in a chem. rxn. is due primarily to energy required to break chem. bonds in reactants & energy produced by forming bonds in products.

15 Enthalpy Balanced eqn. can represent energy absorbed or released during a rxn. C (cr) + O 2(g) CO 2(g) kJ 1 mol C + 1 mol O 2 yield 1 mol CO kj of energy energy released is called Enthalpy of Rxn - ΔH in this case ΔH = kj (exothermic)

16 Standard States Can t meas. enthalpy in absolute terms can only meas. changes in enthalpy depends upon standard of reference used. Our reference will be subst. in their std. states The enthalpy substs. & kPa Std. conditions for thermodynamics are not the same as for gas laws

17 Standard States A change in either temp. or press. can affect enthalpy of a subst. In measuring enthalpy, the enthalpy of free elements are arbitrarily set = to 0

18 Enthalpy of Formation The change in enthalpy when 1 mole of a compound is produced from the free elements in their std. states In: C + O 2 CO 2 + energy C & O 2 have enthalpies of zero Std. molar enthalpy of form. for CO 2 is kj/mol (-) - exothermic rxn. ΔH f is (-) Comps. w/ (+) ΔH f - endothermic

19 Enthalpy of Formation Comps. w/ large (-) enthalpies of formation are thermodynamically stable Thermodynamic Stability - depends on amt. of energy required to decompose the comp. 1 mole CO 2 would need kj to decompose it.

20 Enthalpy of Formation Hg(OCN)2 produces 268 kj when 1 mole decomposes has (+) enthalpy of formation - explosive ΔH f o - enthalpy of formation o indicates values are at kpa & 25 o C f - formation

21 Calculation of Enthalpy of Rxn Enthalpy of products (Σ Δh f o (Products) ) must = enthalpy of reactants (Σ ΔΗ f o (reactants) ) + any change in enthalpy (ΔH ro ) during rxn. Law of Conservation of energy Σ means sum of

22 Calculation of Enthalpy of Rxn Σ ΔH f o (Products) = Σ ΔΗ f o (reactants) + ΔH r o ΔH r o = Σ ΔH f o (Products) - Σ ΔΗ f o (reactants) ΔH f o for free elements is zero Using this, we can predict whether a rxn. will be endothermic or exothermic. If ΔΗ f o is (-), rxn. is exothermic If ΔΗ f o is (+), rxn. is endothermic

23 Hess s Law Some rxns. can be broken down into 1 or more smaller rxns. Hess s Law - The change of enthalpy for a rxn. is the sum of the enthlpy changes for a series or rxns. that add up to the overall rxn.

24 Hess s Law C + O 2 CO 2 ΔH = kj 2 CO CO 2 + C ΔH = kj 2CO + O 2 2CO 2 ΔH = kj Write eqns. so things to be cancelled are on opposite sides of yield sign. Do Sample Prob. #7

25 Entropy Highly exothermic rxns. go spontaneously Weak exo. & endothermic rxns. may also be spontaneous May proceed under strong rxn. conditions ie. temp. incr.

26 Entropy C (cr) + H 2 O (g) + energy Endothermic rxn. - ΔH is (+) CO (g) + H 2(g) If 1 mole C reacts w/ 1 mole H 2 O(g), then ΔH = 131 kj Since for most spontaneous rxns. ΔH is (-), an additional factor must be considered

27 Entropy Entropy - degree of disorder - S Crystalline solids - very orderly arrangement liquids - somewhat less ordered gases - no order high temp. has less order than low temp. Change in Entropy - ΔS - state function

28 Entropy 2 H 2 O 2H 2 + O 2 more particles after rxn. than before products represent a more disordered system

29 Entropy Systems in which the molecs. are far apart are more disordered than those w/ particles close together. 2 liquids dissolved in ea. other make a more disordered syst. than 2 separate liquids. When an obj. is heated, molec. motion incr. & entropy is increased.

30 Entropy (+) ΔS means incr. in degree of disorder syst. becomes less ordered ex) a solid is converted to liquid or gas If opposite occurs (liquid or gas solid) ΔS is (-)

31 Gibbs Free Energy If a rxn. is exothermic, but involves an incr. in order, it may or may not occur. An endothermic rxn. that produces disorder may or may not occur. In both cases, either the enthalpy chg (-ΔH) or entropy change (+ΔS) would make the rxns. spontaneous while the other quantity would prevent the rxns.

32 Gibbs Free Energy Gibbs Free Energy- indicates whether or not a rxn. will occur ΔG - change in Gibbs Free Energy - state function Defined in terms of enthalpy & entropy G = H - TS; ΔG = ΔH - TΔS T is temp. in kelvin

33 Gibbs Free Energy In a spontaneous rxn., ΔG is always (-) Exergonic - spontaneous rxn. ΔG < 0, (-) Endergonic - nonspontaneous rxn. ΔG >0, (+) If a rxn. takes low temp. & has little change in entropy, TΔS is negligible ΔG is mostly a function of ΔH

34 Gibbs Free Energy most spontaneous room temp. have a (-) ΔH Highly endothermic rxns. occur only if TΔS is large. Either if temp. is high or there s a large incr. in entropy.

35 Gibbs Free Energy If ΔH & ΔS have same sign, there s some which ΔH - TΔS will = 0 ΔG = 0 This state is the thermodynamic def. of a system in equilib., G a minimum for the syst.

36 Gibbs Free Energy (Summary) If ΔG is (-), rxn. is spontaneous If ΔG is (+)m rxn. is not spontaneous If ΔG = 0, rxn. equilibrium Changes in nature tend toward low energy state (lg (-) ΔH) and a state of high randomness (lg.(+) ΔS) All spontaneous rxns. proceed towrd equilib. G - free energy- chemical potential energy is least when syst. equilib.

37 Gibbs Free Energy Calculations Enthalpy Change for rxn: ΔH r o = ΣΔH f o (products) - ΣΔH f o (reactants) Free Energy Change for rxn.: ΔG r o = ΣΔG f o (products) - ΣΔG f o (reactants) Entropy Change for rxn.: ΔS o r = ΣSo (products) - ΣSo (reactants) ΔG = ΔH - TΔS

38 Gibbs Free Energy & Equilibrium Gibbs Free Energy is related to K eq by the equation: ΔG = -2.30RT(log K eq ) R = kj/mol K Do #18 & #19 p. 704