Key EquationsChapter 16 - Definition of Pkaop Ka=−log Ka- Henderson-Hasselbalch EquationopH= p Ka+log[Conjugate Base][ Acid]Chapter 17- S = k ln Wo Relating entropy to number of microstates-∆ Suniv = ∆Ssys + ∆Ssurr > 0o Second law of thermodynamics (spontaneous process)- ∆Suniv = ∆Ssys + ∆Ssurr = 0o Second law of thermodynamics (equilibrium process)-∆ Srxn˚ = ΣnS˚(products) – ΣmS˚ (reactants) o Standard entropy change of a reaction- ∆G = ∆H – T∆S o Free-energy change at constant temperatureCHM1046 EXAM 3 STUDY GUIDE- ∆Grxn˚=Σ n ∆ Gf˚(products)−Σ m∆ Gf˚(reactants)o Standard free-energy change of a reaction- ∆G = ∆G˚ + RT lnQo Relationship between free-energy change and standard free-energy change and reaction quotient. - ∆G˚= −RT ln Ko Relationship between standard free-energy change and the equilibrium constant. Chapter 18- E˚Cell = Ecathode˚ - Eanode˚o Calculating the standard emf of a galvanic cell. - ∆G = -nFEcello Relating free-energy change to the emf of the cell.- ∆G˚ = -nFE˚cello Relating the standard free-energy change to the standard emf of the cell. - E˚cell = 0.0257 Vnln Ko Relating the standard emf of the cell to the equilibrium constant. - E˚cell = 0.0592Vnlog Ko Relating the standard emf of the cell to the equilibrium constant. -E = E˚ - 0.0257 Vnln Qo Nernst Equation - Relating the emf of the cell to the concentrations under nonstandard state conditions. -E = E˚ - 0.0592nlog Qo Nernst Equation - Relating the emf of the cell to the concentrations under nonstandard state conditions. Chapter 16 Concepts- Common Ion – The same or common ion that appears in a solution with two dissolved solutes. - Common Ion Effect – The shift in equilibrium caused by the addition of a compound having an ion in common with the dissolved substance. o Plays an important role in all problems that determine the pH of a solution and molar solubility. o The presence of a common ion decreases the solubility of a slightly soluble salt. o Le Châtelier’s Principle = Other name for the common ion effect. - Buffer Solutions – A combination of either a weak acid and its weak conjugate base, or a weak base and its weak conjugate acid.o A buffer solution reacts with small amounts of added acid or base in such a way that the pH of the solution doesn’t change. - pH at Equivalence Points: o Strong acid – Strong base titrations pH at equivalence point = 7 o Weak acid – Strong base titrations pH at equivalence point > 7o Strong Acid – Weak base titrations pH at equivalence point < 7- Acid-Base Indicators – Weak organic acids or bases that change color near theirequivalence point in an acid-base neutralization reaction. o End Point – The point in a titration when the indicator changes color. - Solubility Product (KSP) – Expresses equilibrium between a solid (precipitate) and its ions in solution. o Solubility can be calculated from KSPo KSP can be calculated from Solubility - Molar Solubility – Moles of solute divided by the amount of Liters solution. - Solubility – Grams of solute divided by the amount of liters solution. - The solubility of salts with anions derived from strong acids are unaffected by pH. oAgCl ↔ Ag+¿¿(aq) + Cl- (aq)o Because Cl- is the conjugate base of a strong acid (HCI), the solubility of AgCl isnot affected by an acid solution.- Formation Constant (Kf) – The equilibrium constant for the complex ion formation. o Kf is used to measure the tendency of a metal ion to form a particular complex ion. o Also called stability constant. o Larger Kf = More stable the complex ion- Complex Ions – An ion containing a central metal cation bonded to one or more molecules or ions. o Complex Ions are formed in solution by the combination of a metal Cation with aLewis base. The formation constant Kf measures the tendency toward the formation of a specific complex ion. o Complex ion formation can increase the solubility of an insoluble substance. - Qualitative Analysis – The identification of cations and anions in solution. Chapter 17 Concepts- Entropy – A measure of the different ways a system can disperse its energy. Any spontaneous process must lead to a net increase in entropy in the universe (second law of thermodynamics).- Standard Entropy of Reaction (∆Srxn˚¿- Calculated by the difference in standard entropies between products and rectants. o∆ Srxn˚=Σn S˚(products)−Σ m S˚(reactants)- Standard Free-Energy of Reaction (∆Grxn˚¿ – The free-energy change for a reaction when it occurs under standard-state conditions, when reactants in their standard states are converted to products in their standard states. - Standard Free-Energy of Formation – The free-energy change that occurswhen 1 mole of the compound is synthesized from its elements in their standard states. - Third Law of Thermodynamics – The entropy of a perfect crystalline substance is zero at 0 K. This law enables us to measure the absolute entropies of substances. - Under conditions of constant temperature and pressure, the free-energy change ∆G is lessthan zero for a spontaneous process and greater than zero for a nonspontaneous process. o ∆G < 0 Spontaneous Processo ∆G > 0 Nonspontaneous Processo ∆G = 0 Equilibrium Processo ∆G = ∆H – T∆S Chemical or Physical Process at Constant Temp and Pressure This last equation can be used to predict the spontaneity of a process H, S, and G are state functions.o If ∆H is negative and ∆S is positive, then ∆G will always be negative regardlessof temperature. o If ∆H is negative and ∆S is negative, then ∆G will be negative only when T∆S issmaller in magnitude than ∆H. This condition is met when T is small. - Standard Free-Energy Change for a Reaction (∆G˚)– Can be calculatedfrom the standard free energies of formation of reactants and products. o Free Energy – The energy available to do work. o Gibbs Free Energy = Other name for Free Energy. - The equilibrium constant (k) of a reaction and the standard free-energy change of thereaction (∆G˚) are related by the equation ∆G˚ = –RT ln K- Many biological reactions are nonspontaneous and are driven by the hydrolysis of ATP,for which ∆G˚ is negative. Chapter 18- Redox Reactions – Reactions that involve the transfer of electrons. Any equation representing redox processes can be balanced using the ion-electron method. o All
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