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Chemistry 1200B.docx

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Department
Chemistry
Course
Chemistry 1027A/B
Professor
Kay Calvin
Semester
Fall

Description
Christina Wang Chemistry Study Notes Chemistry 1200B 1.1Equilibrium Constant  Equilibrium reactions: o Reversible processes with equal forward and reverse rates o Many of these reactions look static, but at molecular level there are forward and reverse reaction o When complete will have both products and reactants present, but not in same amounts  Equilibrium occurs when forward reaction rate = reverse reaction rate o i.e. evaporation and condensation of water o With coloured gas (i.e.2N4O and N2 ) the mix will be an in-between colour, once at equilibrium the amount of each gas will not change, thus colour will not change either  Experiment showed a relationship with PRESSURE OF REACTANTS AND PRESSURE OF PRODUCTS o K, is called the equilibrium constant o Kc= measured in concentration units o Kp= measured in pressure units  The value of K will increase/decrease according to temperature fluctuations and if the reaction is endothermic vs. exothermic  Homogeneous equilibrium – all components are in same phase  Heterogeneous equilibrium –not all components are in same phase  SOLIDS ARE NOT INCLUDED IN EQUILIBRIUM CONSTANT EXPRESSIONS o Only gases and liquids  Generally aA + bB ↔ cC + dD K =  If no gaseous or aqueous products are included, write in a “ 1 ”  When reversing equilibrium reactions, equilibrium constant expression must be inverted o K’ = 1/K  When equation in equilibrium is multiplied by two, K expression is squared 2 o K’ = K  When equation in equilibrium is divided by two, K expression is square rooted 1/2 o K’ = K  When combining equilibria, multiply the original K expressions together 1 2 o K’ = K x K  Magnitude of equilibrium constant indicates the relative amount of products or reactant in it o K > 1, equilibrium lies towards products (more products than reactants) o K = 0, equilibrium lies in the middle o K < 1, equilibrium lies towards reactants (more reactants than products)  Reaction quotient – indicates if equation is at equilibrium and if not which direction to proceed o Equation used to calculate reaction quotient Q, is the same as for K o Comparing Q and K, will indicate which direction to proceed to attain equilibrium Page 1 of 18 Christina Wang Chemistry Study Notes o Q = K, equation is at equilibrium o K > Q, reaction proceeds forward (not enough product/too much reactant) o K < Q, reaction proceeds in reverse (not enough reactant/too much product)  A system can be disturbed and it will react to re-establish equilibrium again o Adding/removing reactant/product o Changing pressure/temperature  Value of equilibrium constant changes with the change of temperature o In every other case, only the position of the equilibrium will change (different values, but same overall ratio) FOR: N 2(g) 3H2(g) 2NH 3(g)  Concentration changes (value of K does not change): o If 2 is increased, system attempts to more of the additional2H , reaction proceeds RIGHT  Will use up more N 2s well so there will be an increase in 2N3  Overall: concentration of N is larger, 3H is smaller, and 2NH is larger 2 2 3 o If NH3is decreased, system tries to make more of it, reaction proceeds RIGHT  Decreases amounts of N and H 2 2 o If NH3is increased, opposite will occur, reaction proceeds LEFT  Increases amounts of N 2nd H 2 o All of the above can be explained by looking at comparison between K and Q  If concentration of H2is increased, K is bigger than Q, proceeds RIGHT  If concentration of NH3is increased, K is smaller than Q, proceeds LEFT  Pressure changes (value of K does not change): o If equilibrium pressure of reaction/product is changed result is same as concentration change o Increasing volume, results in drop in pressure, reaction proceeds towards side with more moles o Decreasing volume, pressure will increase, reaction proceeds towards side with fewer moles o If the number of moles of gases are equal on both sides, there is no change in equilibrium o Addition of an inert gas results in an increase in total pressure, but no change in partial pressure thus there is no change in equilibrium position o Systems with only solids/liquids are not affected by change in container volume  Temperature (value of k changes): o + ΔH for a reaction = endothermic; heat is absorbed from surroundings  N 2(g) 3H2(g) 2NH 3(g) 92.0 kJ, produces heat and increases temperature in surroundings o - ΔH for reaction = exothermic; heat is released to the surroundings  N 2(g) 3H2(g) 181 kJ = 2NH 3(g) requires heat and drops the temperature in surroundings o In EXOthermic reactions (more stable to less stable):  increasing temperature has same effect as increasing products  Shift to LEFT, increasing amount of reactants AND decreasing value of K  decreasing temperature has same effect as removing products Page 2 of 18 Christina Wang Chemistry Study Notes  Shift to RIGHT, increasing amount of reactants AND decreasing value of K o In ENDOthermic reactions (less stable to more stable):  Increasing temperature has same effect as increasing reactants  Thus increasing products and equilibrium shifts RIGHT and increasing value of K  Decreasing temperature similar to decreasing reactants  Thus equilibrium shifts LEFT and increases reactants, and decreases value of K  Gibbs free energy determines if reaction is spontaneous or not under standard conditions (25 °C, all gases at 1atm, all species in solution are 1M): o ΔG < 0 reaction proceeds spontaneously, equilibrium lies RIGHT and K is large o ΔG = 0 reaction is at equilibrium (moles of reactant/product won’t change over time) o ΔG > 0 reaction doesn’t occur spontaneously (doesn’t mean reaction doesn’t proceed at all, only proceeds to a small extent), although reverse reaction is spontaneous; lies to LEFT with small K  ΔG = ΔG° + RT ln Q ; under standard conditions ΔG = 0 and Q = K, thus ΔG° = - RT ln K (ΔG° must be in J) -1 -1 o R, the gas constant is 8.314 J mol K  While ΔG° varies with temperature, ΔH° and ΔS° vary little over temperature ranges, thus can use the following equation to find equilibrium constants at temperatures other than 25 °C: Δ ° Δ ° ( ) Δ °  Slope = Δ °  B-intercept =  An equation can be used to determine K at a second temperature once given ΔH° , K and T rxn 1 1 Δ ° ( ) 1.2Solubility of Ionic Compounds  Calcium in solution exists in blood stream, important for nerve transmission, blood pressure regulation, muscle contraction, hormone secretion etc. While the calcium in bones and teeth is insoluble calcium complex; hydroxyapatite o There must be a balance between the soluble one and the insoluble one  Solubility – amount of substance that dissolves in a specific solvent in certain volume at specific temperature o Soluble compound : 10g or more can dissolve in litre of solvent o Slightly soluable: 0.1 - 10g can dissolve in litre of solvent o Insoluble: less than 0.1 g can dissolve in litre of solvent  Solubility guidelines can predict solubility or insolubility Page 3 of 18 Christina Wang Chemistry Study Notes  Solvent – liquid/gas that dissolves another liquid/solid/gas  Solute – substance that dissolves in a solvent  Saturated solution – equilibrium between undissolved solid and dissociated hydrate ions  K is solubility product constant sp  Salts have large Kspnd the dissolutions go till completion y+ x- y+ x x- y  Kspxpression for A Bx y(s) xA (aq)+ yB (aq)is sp= [ A ] [ B ]  Pure solids aren’t included in equilibrium expressions, though solids provide more surface area for dissolution, but reformation also occurs at the surface; doubling SA doubles rate of dissolution but also doubling the rate of precipitation and thus cancelling each other out  Solubility product constant is an equilibrium constant and ALWAYS the same for a given solid at given temperature  Saturated implies at equilibrium between the solid and the ions in solution  Said to be unsaturated if species in solution is less than required for equilibrium  Reaction quotient Q, also known as the ion product is calculated the same way as K sp  For a saturated solution (= at equilibrium) Q = K and can be compared to see if there’s precipitate sp  Ksp Q unsaturated, no solid will be present and ions remain in solution  Ksp Q system at equilibrium (saturated)  K < Q solid will be formed, solution is said to be supersaturated (solution has an ion concentration sp greater than equilibrium values)  Common ion effect can significantly reduce solubility of the solid relative to that in pure water o Reduces amount of solid that will dissolve  AgCl example: Page 4 of 18 Christina Wang Chemistry Study Notes 1.3Weak Acids and Bases  Acids increase hydrogen ion concentration  Bases increase hydroxide ion concentration  Strong acids/bases completely ionize in solution; weak acids/bases do not ionize completely o Weak acid/base reactions proceed until equilibrium  HCl is a strong acid, dissolved in water produces 1 mol H and 1 mol Cl  Strong acids: HBr, HI, HNO3, HClO4, H2SO4  Strong bases: KOH, Mg(OH) ,2Ca(OH) 2  Arrhenius theory: + o Acid produces H ions in water o Base produces OH ions in water  Brownsted-Lowry theory: o Acids are proton donors o Bases are proton acceptors  Weak acid (HA) donates a hydrogen ion to water as it ionizes  Weak base (B) accepts a hydrogen ion as it ionizes  K – equilibrium constant for acid a  Kb– equilibrium constant for base  Lewis acids: accept pair of electrons result in formation of coordinate covalent bond  Lewis base: donates pair of electrons to form an acid-base adduct  Organic amines (ammonia with one or more H atoms substituted by an organic group) behave as Bronsted-Lowry bases o Nitrogen in all amines carries a nonbonding e- that attracts positive hydrogen in water molecules; the hydrogen is transferred to the amine and it donates the nonbonding e- pair to form N-H bond o The nonbonding pair of e- is a required feature of Lewis bases (ammonia donates nonbonding e- pair and hydrogen ion from water accepts them)  Metal cations (coordination complexes) act as Lewis acids; complex ions react with Lewis bases (i.e. water) and forms coordinate bonds  Oxides of non-metals behave as Lewis acids when react with hydroxide  Equilibrium constant expression, K, governs weak acid and weak base equilibrium reactions  Weak acids and weak bases have K values much less than 1; larger K = greater ionization  Increase in K = stronger weak acid/base = decrease in pKaor pK balue Page 5 of 18 Christina Wang Chemistry Study Notes  pH = -log [ H ] pOH = -log [ OH ] pH + pOH = 14  Initial concentration of acid is represented by c  Amount of acid ionized, concentration of hydrogen ions, and conjugate base are represented by x  There is little difference between quadratic way and the approximation way (error is less than 1%)  When to use the approximation solution (if you know x then use it): 1. If % ionization is less than 5% 2. If c/k is more than 400  More diluted solution of weak acid/base = greater % ionization = increase in pH o K doesn’t change, concentration of products at equilibrium must change when initial concentrations of acid/base is decreased by dilution  Conjugate base – one less hydrogen atom than the weak acid counterpart; more negatively charged than weak acid o Behaves like a weak base in solution  Conjugate acid – one more hydrogen than the weak base counterpart; more positively charged than the weak base o Behaves like an acid in solution 14  Kwis 1.0 x 10 = equilibrium constant for ionization of water at 25°C; ion product of water  Kw= K a K bnd pK +apK = b4  Stronger the weak acid = weaker the conjugate base; strong acids have extremely weak conjugate bases o Conjugate bases are so weak they don’t affect pH at all; spectator ions  Salts – ionic solid containing cations and anions o When dissolved in water cations and anions separate; what is dissolved completely ionizes  Spectator ions – cations/anions that don’t effect pH  Some cations (i.e. ammonia’s conjugate acid) act as weak acids making the solution acidic Page 6 of 18 Christina Wang Chemistry Study Notes  Anions that are conjugate bases of weak acids act as weak bases ;sodium acetate ionizes water  K a K b solution is acidic  Amphiprotic anion – species that can acts as both an acid or base, the above statement is used to determine if acts as a base or as an acid, i.e. ammonia and water  ΔG° = - RT ln K o For a weak acid K < 1; ln K = negative; ΔG° =positive number o KHZ K HYus ΔG° < ΔHZ HY o Stronger weak acid = more ionization = more favourable - -  Since HZ is stronger weak acid; Z is more favourable than Y from HY  High in energy = highly localized charge = thermodynamically unstable = less favourable  Relative strength of weak acid is primarily determined by the relative thermodynamic stabilities of the conjugate bases; generally unnecessary to consider unionized weak acids, as they are neutral and their energy differences can usually be ignored o HZ is a stronger acid than HY, because Z is more stable (less in energy) than Y-  What affects conjugate base stability: resonance, electronegativity, and induction 1. Resonance: delocalizes charges over two or more atoms and INCREASES stability of the ion a. Increase in resonance structure = increase delocalization = increase stability = stronger acid b. Effect is cumulative; more resonance structures = stronger acid c. Oxoacids – acids that contain a non-metal atom bonded to 1 or more oxygen atoms; at least one oxygen is bonded to a hydrogen atom 2. Electronegativity: charges are high in energy = unstable species a. Which can better accommodate a negative charge = more electronegative atom b. Every molecule with hydrogen atom can (in principle) donate a proton c. High EN = stronger weak acid d. Electronegativity doesn’t work when going down a group; comparison only works on atoms of similar size 3. Inductive effect: also caused by electronegativity; examines a distance effect a. Results in one or more EN atoms stabilizing a negative charge elsewhere in the ion b. Increased induction = decreased e- density = increased stability = stronger weak acid c. An EN atom pulls the electrons away and further stabilizes causing less electron density d. Also cumulative, more EN atoms present = greater stabilization of conjugate base e. Inductive effect falls with distance, if EN atom is far its effect is weaker  Common ion effect – example of Le Chatelier’s principle Page 7 of 18 Christina Wang Chemistry Study Notes  For example:  Polyprotic weak acids – contain more than one ionisable hydrogen atom o Will cause ionization in stages with a different K value in each stage o Example of diprotic acid H 2O (3imultaneous equilibria; both are happening at the same time) - o Since K a1 larger than K ta2s H CO2is 3onized more than HCO 3 o Hydronium ion is produced in steps, but pH of the solution is mostly due to the production in first step o Ka2s a lot smaller because there are two negative charges = less stability Page 8 of 18 Christina Wang Chemistry Study Notes 1.4 Buffer Solutions  HA + H O ↔ H O + + A- is acidic due to H O + (aq) 2 (l) 3 (aq) (aq) 3  A (aq) H 2 ↔(l) (aq)+ OH -(aq)s basic due to hydrolysis -  Buffer solution – inclusion of both conjugate base(A) and weak acid (HA) o Able to withstand changes in pH when small amounts of strong acid/base is added to it o Requirements: 1. Buffer solution contains weak acid that will react with any OH ions from strong base + 2. Contain a weak base that will react with any H O ions3from strong acid 3. Acid and base don’t react with each other: weak acid/base and its conjugate partner; must
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