Acids and Bases

14.1: The Nature of Acids and Bases Arrhenius postulates that acids produce H+ in solution Bronsted-Lowry: Acid is proton donor, and a base is a proton acceptor Hydronium Ion: H3O+ General form: HA(aq) + H2O(l) → H3O+(aq) + A-((aq) Conjugate base: Everything from acid after proton is lost Conjugate acid: formed when the proton is added to a base Conjugate acid-base pair: Consists of 2 substances related together via donation and acceptance of a single H+ Ka = [H+][A-] / [HA] A- = conjugate base

14.2 Acid Strength Strength of an acid is defined by equilibrium position of its ionization reaction Strong acid: far to the right (almost all HA ionizes) Yields weak conjugate base; weak acids are the opposite Diprotic acids have two acidic protons (i.e. H2SO4) Oxyacids: Acidic hydrogen is attached to oxygen atom\ Organic acids: have carbon atom backbone and usually a carboxyl group (COOH) Some acids are so strong that Ka cannot be calculated Amphoteric = acid and base (i.e. water) Autoionization of water: transfer of proton from one water to another Kw = ion-product constant refers to ionization of water Kw = [H+][OH-] = 1.0 x 10-14

14.3 The pH Scale pH measures concentration of H+ in solution pH = -log[H+] Sig figs: Decimal places in log = sig figs in original number pOH = -log[OH-] pK = -logK pH changed by 10 for every power of 10 change in [H+]

14.4 Calculating the pH of Strong Acid Solutions Major species: solution components present in relatively large amounts Must deal with major species: In 1.0 M HCl, there is essentially no HCl Mostly H+, Cl-, and H2O Use Le Chatelier's principle to determine direction of shift and sources of ions

14.5 Calculating the pH of Weak Acid Solutions  

14.5 Calculating the pH of Weak Acid Solutions  

14.6 Bases -Complete dissociation = strong base, and all 1A hydroxides are strong bases -Finding the pH of strong bases -Find Concentration of [H+] from K: ([OH-] is equal to NaOH because it is such a strong acid) pH = -log(2.0 x 10-13 M) = 12.70 --Even bases without OH- produce OH- via reaction with H2O (lone pairs meet up w/ H in H2O; lone pair normally on Nitrogen) -- Base reaction and water = B(aq) + H2O(l) ---> BH+(aq) + OH-(aq) KB = reaction of base with water, forms conjugate base OH- ---- weak bases

14.7 - Polyprotic acids ---Polyprotic acids can donate more than one proton to solution; triprotic donates 3 ---Conjugate base of first dissociation becomes acid of second (i.e. H2CO3 dissociates to form HCO3- , which becomes CO32-) ---Generally, Ka1 > Ka2 > Ka3…

Polyprotic practice problems pH of H3PO4 and the concentrations of each of its subsequent conjugate bases (all Ka values used here are given in the problem , would take a while to write out all at once)

Sulfuric Acid This acid is unique -- strong acid for first dissociation, weak for the rest I will do this example on the board if necessary As a general rule, if H2SO4is relatively concentrated (think 1.0 M or higher), the usual approximations for monoprotic strong acids will suffice. The second dissociation becomes significant only in dilute concentrations

14.8 -- Acid-Base Properties of Salts Salts are ions Conjugate base of strong acid has little affinity or protons in water. So, salts with anions from strong acids and cations from bases have no effect on pH Ka * Kb = Ww = 1.0 x 10-14 for any weak acid and its conjugate base Any salt with a neutral cation and an anion that is the conjugate base of a weak acid has a basic solution SO: Calculate pH of .3 M NaF, given Ka = 7.2 x 10-4 Use above relationships: Kb = 1.0 x 10-14 / 7.2 x 10-4. ICE it up (w/ approximations) to get: (x^2 / .3) solve for x to find [OH-] and find pOH (ends up being about 5.69) pH = 14 - 5.69 = 8.31

But can Salts make Acidic solutions too? OF COURSE!!!!!!!!!!!!!!!!! Salts with a non-base anion and cation that is the conjugate acid of a weak solutions make acids Solving these problems is the same as any other weak acid problem → you might need to be able to go from Kb to Ka --- remember the relationship from the previous slide Best tip is to just write out dissociation reaction for major species, than find Ka (if not given), than set up your ICE chart (if necessary) Some cations that are highly charged can can act as acids High charge polarizes OH when the the molecules are hydrated, making the hydrogens more acidic than in normal water these problems are the exact same as other weak acids. I can do an example if it is deemed necessary

Structure and its effect on Acid-Base Properties If there’s a hydrogen, it can be an acid, but not always--------why????? Strength and polarity of the bond are absolutely crucial Acids require high polarity OR weak strength or maybe even BOTH In Oxyacids (X--O--H), more O = more acidic If X is electronegative, it also increases acidity -- the more you know!

Acid-Base Properties of Oxides But X-O-H can produce hydroxides and be a base, how will we know which it does? More electronegative X makes the O-X bond less likely to break, meaning that only a proton can be broken off -- higher electronegativity leads to acid, lower electronegativity leads to base

The Lewis Acid-Base Model Lewis always makes it about electron pairs Lewis acid is an electron pair acceptor; a base is an electron pair donor This ecompasses all Bronstead-Lowry acids and than some, so it is very useful

A closing quote, courtesy of our beloved Zumdahls: “Let the problem guide you. Be patient.”