Lecture 13 Announcements (2024, Chemistry)

Lecture #13, p. 1 Lecture 13: Announcements Today: Brown Ch. 16 Acid–Base Equilibria 16.1 Acid–Base Equilibria 16.2 The Autoionization of Water 16.3 The pH Scale...

Lecture #13, p. 1 Lecture 13: Announcements Today: Brown Ch. 16 Acid–Base Equilibria 16.1 Acid–Base Equilibria 16.2 The Autoionization of Water 16.3 The pH Scale 16.4 Strong Acids and Bases 16.5 Weak Acids 16.6 Weak Bases 17.1 The Common-Ion Effect 17.2 Buffers Chemistry  Lecture #13, p. 2 Lecture 13: Announcements Problem Set 12: Due before Exercise #13 tomorrow; upload on Moodle link Exercise #13: Last exercise session! Problem Set 13: Posted on Moodle; due by Friday, Dec. 22, 14:00 Problem Set 14: Posted on Moodle with solutions Do not need to hand in! Study Center: Next Wednesday, 18:00–20:00 in ETA F 5 (last one!) Office Hours: No office hours today Chemistry  Lecture #13, p. 3 Lecture 14 Next Week: Brown Ch. 20 Electrochemistry 20.1 Oxidation States and Oxidation–Reduction Reactions 20.2 Balancing Redox Equations 20.3 Voltaic Cells 20.4 Cell Potentials under Standard Conditions 20.5 Free Energy and Redox Reactions Chemistry  Red Thread Last four weeks? Acid-Base Catalysis Properties Christmas! Kinetics Batteries Equilibrium Chemistry  Lecture #13, p. 4 Review In Lecture 12, we discussed chemical equilibria Chemical rxns experience dynamic equilibria between reactants and product At equilibrium, the forward and reverse reaction rates balance Quantified by equilibrium constant, !! or !" Equilibrium constant given by Law of Mass Action Equilibrium constants are unitless Reaction quotient, ": equilibrium-constant expression away from equilibrium Meaning of ", math with !’s, solubility product, !#" Le Châtelier’s Principle: response to changes in concentrations, #, and $ Catalysts and equilibria: equilibrium achieved faster, but ! is same Relationship between ∆' and ! Chemistry  Lecture #13, p. 5 Limescale (Kalk) Revisited Lecture 3: Precipitation reaction: Ca$% #$ + CO$' & #$ → CaCO& ()) Neutralization reaction: Add vinegar to clean londoncityplumbers.co.uk CaCO& ) + 2CH& COOH #$ → H$ O. + Ca CH& COO $ #$ + CO$ (/) Why does scale deposit more in tea kettle? CaCO& is more soluble in H$ O at higher T Counterintuitive Can we understand? which.co.uk Chemistry  Lecture #13, p. 6 Limescale (Kalk) Revisited Why does scale deposit more in tea kettle? Apply our knowledge to understand... Precipitation reaction is more complicated: Ca%& #$ + 2HCO( ' #$ ⇌ Ca %& #$ + CO%( #$ + CO #$ + H O * ' % % Ca%& #$ + CO%( ' #$ ⇌ CaCO' (-) Why does heating water lead to scale deposits? Chemistry  Lecture #13, p. 7 Today: Acid–Base Equilibria Lecture 3: Acid ≡ “substance that ionizes in H2O to form protons” + − Ex: HCl(&') → H (&') + Cl (&') (hydrochloric acid, strong acid) Base ≡ “substance that accepts (reacts with) protons” + − Ex: NaOH &' → Na (&') + OH (&') (sodium hydroxide, strong base) But promised to return: Acid–base chemistry important for engineers to understand Now that we have learned about equilibria, we can go further with acid–base We can be more precise about what they are and how they behave Today: important topic! Chemistry  Lecture #13, p. 8 Definition of Acids and Bases Not so easy... Arrhenius: Acid ≡ “a substance that, when dissolved in water, increases [H+]” !! " + − Ex: HCl(&') H (&') + Cl (&') Base ≡ “a substance that, when dissolved in water, increases [OH−]” !! " + − Ex: NaOH &' Na (&') + OH (&') Useful, but limited: Restricted to aqueous solutions Water is clearly important for acid–base chemistry We will focus on aqueous acid–base, but concept is broader Other definitions? Chemistry  Lecture #13, p. 9 Toaddressthisquestion let'sfirstdiscussprotons Htandwhathappensto Ht inH2o H+ in H2O ? H+ is bare proton: Small, positive, and seeking electrons What happens to proton placed in water ? hydronium ion H+ reacts with lone pair on H2O Forms hydronium ion: H3O+ H3O+ then forms more complex structures with H2O H+ and H3O+ often used interchangeably, but H3O+ more accurate visiting For Chemistry exg Thatis indiscussionsof aqueous acidbasechemistry H H30ᵗ  Lecture #13, p. 10 Wecanthendiscussthecentralequilibriumin aqueous acidbasechemistry Autoionization of Water Much of acid–base chemistry is in H2O Iiiiii Key equilibrium: H2O (ℓ) + H2O (ℓ) ⇌ OH− (aq) + H3O+ (aq) A proton transfers between H2O molecules Elastin ion-product "! = [OH−] [H3O+] ≡ "" ≡ constant = 1.0 × 10−14 (25 °C) iiiiii This equilibrium constant is so important, it is given a special name Note: "" ≪ 1 ! Only 2 of 109 water molecules are ionized at a time ∴ Pure water is almost entirely H2O 25 Chemistry 20 le Htaq OH aq autoionizationreaction Alternativeforms with Ht instead of Hot Kw Ht OH 1.0 1014 25C ionproduct constant  Lecture #13, p. 11 So if thereis solittleHot whythenis thisimportant Why Do We Care? !! = [OH−] [H3O+] When [H3O+] = [OH−] ⇒ solution is “neutral” iii l But often [H3O+] ≠ [OH−] ⇒ If [H3O+] increases, [OH−] decreases iiiIii If [H3O+] > [OH−] ⇒ “acidic” If [H3O+] < [OH−] ⇒ “basic” Moreover, '# applies to any aqueous solution ! We ignore affect of other ions on the equilibrium Thus, '# = [OH−] [H3O+] can be used to manipulate [H3O+] We can change [H3O+] over many orders of magnitude ! Chemistry  Lecture #13, p. 12 Thisleadstoourmeasurementsystemfor 4301 How Do We Quantify? Log scale suited to quantities that change by orders of magnitude pH = −log H"O# !! = [OH−] [H3O+] = 1.0 ×10−14 (25 °C) H"O# = 1.0 × 10−7 pH = −log 1.0 × 10−7 Iiiiiiii Neutral H2O OH$ = 1.0 × 10−7 = 7.00 1 H"O# = 1.0 × 10−3 pH = 3.00 If we increase [H3 O+ ] by 104 OH$ = 1.0 × 10−11 Ti Acidic: pH < 7 and Basic: pH > 7 Chemistry Alsoknownasalkaline  Lecture #13, p. 13 SospecificallyhowdowemanipulatetheautoionizationofH2o How Do We Manipulate? '# = [OH−] [H3O+] Acids: transfer H+ to H2O to make more H3O+ Ex: HCl + + H$ O ℓ Cl& /0 + H' O( (/0) Bases: accept H+ to H2O to make more OH− Ex: NH' /0 + H$ O ℓ NH)( /0 + OH & (/0) i.e., proton- Plus, this approach to acid–base applies beyond H2O transfer approach Ex: HCl + + NH' /0 NH' Cl 4 This is a gas-phase reaction that does not involve water Chemistry  Lecture #13, p. 14 Nowwecanintroduceabroaderdefinition ofacidsandbases Brønsted-Lowry Acids and Bases Brønsted-Lowry Acid: A substance (molecule or ion) that donates a proton to another substance Brønsted-Lowry Base: A substance (molecule or ion) that accepts a proton from another substance Broader definition: water must not be involved Chemistry  Lecture #13, p. 15 Acid–Base Pairs Proton transfer requires acid–base pair Amphiprotic substance: Autoionization of H2O ⇒ H2O plays both roles ⇒ can act as both H2O (ℓ) + H2O (ℓ) ⇌ OH− (aq) + H3O+ (aq) acid or base We can write generic reaction of acid “HA” with H2O: HA (aq) + H2O (ℓ) ⇌ A− (aq) + H3O+ (aq) "- = A' [H& O% ] 1 iii [HA] acid base conjugate conjugate acid-dissociation base acid constant H+ H+ donor acceptor Chemistry  Lecture #13, p. 16 Acid–Base Pairs We can write generic reaction of base “B” with H2O: HB% [OH ' ] B (aq) + H2O (ℓ) ⇌ HB+ (aq) + OH− (aq) ". = [B] base acid conjugate conjugate base-dissociation acid base constant H+ H+ acceptor donor true For both reactions ⇒ LHS: acid + base RHS: conjugate acid + conjugate base s h Each acid, HA has conjugate base, A− s base, B conjugate acid, HB+ site Chemistry  Lecture #13, p. 17 Strength of Acids HA (aq) + H2O (ℓ) ⇌ A− (aq) + H3O+ (aq) acid base conjugate conjugate base acid Strong acid, HA: Gives up all H+; HA is fully ionized HA completely transfers H+ to H2O to form H3O+ Conjugate base A− does not accept protons Decreasing acidity Conjugate base has negligible basicity Weak acid, HA: Partially ionized; mix of HA and A− Conjugate base is also weak because it also only partially accepts protons Weak conjugate base A− HA without acidity: Gives up no H+; A− is strong H+ acceptor Strong conjugate base A− Chemistry  Lecture #13, p. 18 Strength of Bases B (aq) + H2O (ℓ) ⇌ HB+ (aq) + OH− (aq) Similarly for bases base acid conjugate conjugate acid base Strong base, B: Grabs all H+; HB+ dominates Decreasing basicity Conjugate acid HB+ does not donate protons Negligible acidity Weak base, B: Grabs some H+; mix of B and HB+ Weak conjugate acid HB+ B without basicity: Grabs no H+; HB+ is strong H+ donor Strong conjugate acid HB+ Chemistry  Lecture #13, p. 19 Strength of Acids and Bases Chemistry  Lecture #13, p. 20 Strength of Acids and Bases Strong acids: Strong “givers”: Give all H+ to H2O Not present as HA in H2O, rather as A− Strong bases: Strong “takers”: Take H+ from H2O Not present as B in H2O, rather as HB+ Explains (Brown, p. 762): E.IE ii “The ions H3O+ and OH− are, respectively, the strongest possible acid and strongest possible base that can exist at equilibrium in aqueous solution.” Chemistry  Lecture #13, p. 21 Strength of Acids and Bases Explains (Brown, p. 762): “The ions H3O+ and OH− are, respectively, the strongest possible acid and strongest possible base that can exist at equilibrium in aqueous solution.” “Stronger acids react with water to produce H3O+ ions, and stronger bases react with water to produce OH − ions, a phenomenon known as the leveling effect.” These H3O+ and OH − ions are stronger acids and bases than most other chemicals. Thus, can give/take H+. Reactive! That’s why we care! Chemistry Even if itof OH 1 107 Small Butmakes a difference Ex Ia as t.is efHY  Lecture #13, p. 22 pH Scale and Other “p” Scales p{} = −log {} in general From above: H' O 5 + H' O 5 ⇌ H" O# 78 + OH $ (78) !! = H" O# OH $ = 1.0×10$%& pH = −log [H" O# ] Other “p” scales: pOH = −log [OH $ ] p!! = −log !! = 14.00 = −log{ H" O# OH $ } = pH + pOH = 14.00 Chemistry  Lecture #13, p. 23 pH and pOH Scales Chemistry  Lecture #13, p. 24 pH Scale and Other “p” Scales p{} = −log {} in general From above: H$ O 1 + H$ O 1 ⇌ H& O% 45 + OH ' (45) !* = H& O% OH ' = 1.0×10'/, pH = −log [H& O% ] Other “p” scales: pOH = −log [OH ' ] p!* = −log !* = 14.00 = −log{ H& O% OH ' } = pH + pOH = 14.00 0" (# )$ HA #$ + H% O ( ⇌ A& #$ + H' O( #$ !- = [(0] p!- = −log [!- ] The larger Ka, the stronger the acid. The smaller pKa, the stronger the acid. (3$ )(" B #$ + H% O ( ⇌ HB ( #$ + OH & #$ !. = p!. = −log [!. ] The larger Kb, the stronger the base. The smaller pKb, the stronger the base. Chemistry  Lecture #13, p. 25 Values of Ka and Kb? B #$ + H% O ( ⇌ HB ( #$ + OH & #$ !. = (3$ )(" HA #$ + H% O ( ⇌ A& #$ + H' O( #$ 0" (# )$ !- = [(0] Chemistry  Lecture #13, p. 26 Meaning of pKa? Comes up a lot in chemistry, biology, and materials science... Has direct physical meaning: 0" (# )$ HA #$ + H% O ( ⇌ A& #$ + H' O( #$ !- = p!- = −log [!- ] [(0] A& H' O( A& H' O( When: pH = p-) −log H' O( = −log [HA] H' O ( = [HA] [HA] = A' ! Thus: when pH = p!- Acid is half ionized when pH < p!- Acid is mostly protonated: exists as HA when pH > p!- Acid is mostly charged: exists as A− Chemistry 1Important forexample if wehave a surfacee.g glass withacidicgroupson it Is thesurfacecharged at agivenpH  Lecture #13, p. 27 Common-Ion Effect Add common ion to manipulate acid-base equilibria Ex: CH& COOH $% + H$ O ℓ ⇌ CH& COO' $% + H& O% ($%) acetic acid acetate ion Add strong electrolyte: CH& COONa completely CH& COONa $% + H$ O ℓ CH& COO' $% + Na% ($%) dissociates sodium acetate acetate ion Increase in acetate ion shifts acid-base equilibrium left Le Châtelier’s Decreases H& O% and increases pH Principle In words: Whenever a weak electrolyte and a strong electrolyte containing a common ion are together in solution, the weak electrolyte ionizes less than if it were alone in solution Chemistry Note Equilibriumconstantdoesnotchange Justrelativeconcentrations ofproductsand reactants Canalsoworkwithweakbaseloweringoh  Lecture #13, p. 28 Whocares Veryuseful forcontrollingpHofbuffers Sowhatis abuffer Buffers Common-ion effect is very useful for controlling pH of buffers Many industrial and biological processes are sensitive to pH changes Buffers protect against pH disturbances and maintain pH Fiange How? Weak acid–base conjugate pair protects against added [H3O+] or [OH−] Disturbance HA (aq) + OH− (aq) ⇌ A− (aq) + H2O (ℓ) [HA] Sets [H& O% ] = 1- A− (aq) + H3O+ (aq) ⇌ HA (aq) + H2O (ℓ) [A' ] pH an As long as disturbance is small relative to [HA], [A' ] ⇒ small pH change Chemistry  Lecture #13, p. 29 Notes We can tune the buffer pH with common-ion effect Add MA to increase [A−] Ex: CH! COOH $% + CH! COONa $% [HA] [MA] ⇒ increases [A−] [base] Equilibrium To calculate buffer pH use: pH = p," + log concentrations [acid] [A−], [HA] Henderson-Hasselbalch Equation To effectively resist pH changes in both directions: [HA] = [A# ] pH pka It is possible to overwhelm buffer if disturbance is too large Blood is an effective buffer using H2CO3, HCO3− pair Chemistry carbonifacid bifarbonate  Lecture #13, p. 30 Carbonate System 1. CO! + H!O ⇌ H!CO" 2. H!CO" + H!O ⇌ HCO#$ " + H"O % 3. HCO#$ #! " + H!O ⇌ CO" + H"O % Blood Buffer If [CO2] in blood rises, equilibria of 1 and 2 shift right, and pH lowers Lungs lower [CO2] in blood, equilibria shift left, and pH rises If pH is too high, kidneys remove HCO3− from blood, and pH lowers Also plays an important role in the ocean... Chemistry  Lecture #13, p. 31 Engineers Can Use This! “Direct Ocean Capture” Captura Next time: Batteries!  Lecture #13, p. 32 What We Learned Definitions of acids and bases Autoionization of water, ion-product constant, 1* Proton-transfer reactions, Brønsted-Lowry Acids and Bases Acid–base pairs, conjugate acids and conjugate bases Acid-dissociation constant, 1- Base-dissociation constant, 1. pH, pOH, p1- , p1. , p1* Common-ion effect Buffers, Henderson–Hasselbalch Equation Chemistry