Overview Guide Chp. 19 – Entropy and Free Energy Chp. 17 – Acids, Bases, Equilibrium Chp. 18 – Buffer solutions, pH, Effects of Common Ions Chp. 19 Entropy and Free Energy Reaction Spontaneity – rxns are said to be spontaneous when they occur on their own without outside interference; they "go" on their own. Rxns that are non-spontaneous do not "go" on their own, and need intervention (usually includes putting energy into the reaction in the form of heat). Entropy – is the dispersal of energy throughout a number of matter particles. Second Law of Thermodynamics – energy will always disperse (spread out), unless prevented from doing so. This is because the probability that the energy will spread to a number of particles is extraordinarily higher than the probability that energy will remain concentrated in a few particles. Think of it as a "diffusion" of energy to particles. Third Law of Thermodynamics – there is no disorder in a perfect crystal at 0K. T he result is that ∆ S = q rev /T. q is the heat absorbed, and T is the Kelvin temperature at which the change occurs. Standard Entropy S º – entropy gained by from perfect crystal to standard state conditions, 1 bar, 1 molar. Comparing Entropies: – Entropies o f Gases are much higher than Liquids – Liquids are higher than solids – Larger and more complex molecules have larger entropies than smaller simpler molecules – Dissolved Gases have lower entropy than non-dissolved gases – When solids or liquids are dissolved in a solvent, their entropy increases. – Weaker ions have larger entropies; stronger ions have smaller entropies: example: NaCl(s)>MgO(s) **with these general rules, you should be able to predict whether a reaction h as a negative or positive change in entropy. Gibbs Free Energy – ∆ G º = ∆ H º -T ∆ S – The most important thing to remember is that a negative ∆ G is a spontaneous process; a positive ∆ G is a non-spontaneous process. – negative ∆ H is called an exothermic process ; a positive ∆ H is called an endothermic process – Some reactions can either be spontaneous or non-spontaneous depending on temperature – Some reactions are only spontaneous and some are only non-spontaneous. Example: if ∆ H is positive and ∆ S is negative, the rxn will always be non-spontaneous and unreversible, since ∆ G is always negative. Experiment with this and convince yourself of the various cases. **Remember how to calculate ∆ G º , moles of product times ∆ G º (product) – moles of reactants times ∆ G º (rea ctants). If it is not in standard conditions, ∆ G of products and reactants should be calculated using the ∆ H º , T, and ∆ S º . Equilibrium and Gibbs Free Energy: ∆ G º = -RTlnK º where K is equilibrium constant, R is 8.314 J/(molxK), and T is temp. in K. Nonsta ndard conditions for Gibbs Free Energy: ∆ G = ∆ G º + RTlnQ Used when pressures and concentrations are not standard (not at 1bar) Ch. 17 Acids, Bases, and Equilibrium Br ø nsted-Lowry acids/bases : acids are H+ donors; bases are H+ acceptors. Conjugate Acids/ Bases: Conjugate acid is the acid resulting from the protonation of a base; conjugate base is the base resulting from the deprotonation of the acid. HA + B- – – – – > A- + HB acid + base – – – > conjugate base + conjugate acid *Acids from H30+ in water **Bases for OH- in water ****It is very important that you memorize the relative strengths of acids/bases and their conjugates. These must, for the most part, be memorized. There are a few rules that may help you: – The more st able the conjugate base is as free floating ion, the stronger the acid. It can easily dissociate. – The less stable the base is as an ion, the stronger the base (it doesn't "like" to be in its ion form, so it steals H+ to become more stable). Auto-ioniz ation of Water: Water naturally exists as a mixture of H30+, H20, and OH-. Kw = [OH-][H+] = 10^-14 pKw = -log(Kw) = 14 neutral: pH = 7, [H+] = [OH-] = 10^-7 acidic solutions: [H+] > 10^-7 and [OH-] < 10^-7, pH < 7 basic solutions: [OH-] > 10^ -7 and [H+] < 10^-7, pH > 7 An acid is strong when it completely dissociates and [OH-] is negligible, pH < ~2 A base is strong when [H+] is negligible, pH > ~12 HCl is a strong acid (because it dissociate completely). CH3- is a very strong base. Carboxylic Acids: RCOOH, moderately acidic Amines: RNH2, moderately basic Lewis Acids/Bases : Bases are electron donors; acids are electron acceptors. example: positive metal ions as lewis acids Al(OH)3(s) + OH – – – > Al(OH)4- Chp. 18 Buffer Solutions, Common Ion Effects Buffer Solution – a solution of about equal amounts of weak acid and its conjugate base (or weak base and its conjugate acid). Adding base or acid will not change pH of solution very much. Buffers are retain thei r buffering capability to about +/- 1 pH unit. pH of buffer solutions – – Henderson-Hasselbalch equation pH = pKa + log[conjugate base]/[conjugate acid] **equation is only valid for +/- one pH unit relative to pKa, thus [conj base]/[conj acid] ratio must b e 0.1 and 10. – adding acid or base to a buffer solution changes the [conj base]/[conj acid] ratio by [conj base -/+ x]/[conj acid +/- x] Titrations – adding acid or base to a solution to record change in pH. – endpoint is where buffer of solution breaks d own (pH rises rapidly) – at halfway to endpoint, pH is equal to the pKa. Solubility and Ksp – for AB(s) – – > A + B Ksp = [A][B]; the higher Ksp, the more soluble the solid is. example: AgCl(s) – – – > Ag+ + Cl- Ksp = [6.3 X 10^-6][6.3 X 10^-6] = 4.0 X 10^-11 Common ion effect – presence of COMMON ions reduces solubility (Le Chatler's Principle).