Acids n’ Bases

1. Acids n’ Bases By Jenna Kainic & Emma Martin

2. Types of Reactions • Neutralization: Acid + base  salt + water ex. HCl + NaOH  NaCl + H2O • Gas-Forming: A base without an actual OH- group reacts with H+ to form a molecular compound, such as H2S (g). ex. 2HCl + Na2S  H2S + 2NaCl

3. Definitions • Arrhenius: An acid increases [H+], while a base increases [OH-]. • Brønsted-Lowry: An acid is a proton donor, and a base is a proton acceptor. • Lewis: An acid is an electron-pair acceptor, and a base is an electron-pair donor.

4. Autoionization of Water • At any given time, a very small number of water molecules act as proton donors (acids) to other water molecules, which accept them (bases). • H2O  H+ + OH- ; 2H2O  H3O+ + OH- • Because this is possible, water has an equilibrium constant, which is useful. [H+][OH-] = 1.0 × 10-14 = Kw(at 25ºC)

5. Strong Acids/Bases • Completely dissociate when placed in water into H+ and an anion or a cation and OH-, respectively. • No K value because they are not in equilibrium. The reaction proceeds only in the forward direction. • Therefore, the concentration of the acid is also the concentration of H+ions.

6. Weak Acids/Bases • Do not completely dissociate in water and are therefore in equilibrium. • Ka/Kb is the equilibrium constant. • Ka=[H+][A-]/[HA] • Kb=[B+][OH-]/[BOH] • pKa/b= -log(Ka/b) • pH= -log([H+]) • (Ka)(Kb)=[H+][OH-]=Kwand that’s why it’s useful, folks.

7. Polyprotic Acids • Have multiple H+ ions per molecule • Dissocation occurs in multiple steps • Different equilibrium constants: Ka1, Ka2, Ka3 • Which one do you use? It becomes very obvious on a titration curve.

8. Titration Curves

9. Visual Help with Titrations

10. Titrations: Now It’s Your Turn • A 20.0 mL sample of 0.250 H2SO3solution is titrated with 0.250 M NaOH solution. • Ka1=1.7×10-2 ; Ka2=6.4×10-8 • Calculate the pH when the amount of added NaOH solution is 0, 10, 20, 25, 30, 40, 50, and 55. • We WILL call on all of you, so wake up.

11. Common-Ion Effect • If a strong electrolyte and weak electrolyte in solution contain a common ion (ex NaC2H3O2 and HC2H3O2), the weak electrolyte will be ionized less. • This follows Le Châtelier’s Principle. • This effect is very useful in creating buffers/explains how buffers work.

12. Buffers!!!!1! • Composed of a weak acid and its conjugate base (following the last example, C2H3O2-and HC2H3O2). • Capable of absorbing large changes in pH. This capability differs from buffer to buffer and is called buffer capacity. • Henderson-Hasselbach Equation:pH = pKa+ log([base]/[acid])  usually just used for buffers

13. Acidic/Basic Salts • An anion that is the conjugate base of a strong acid (ex. Cl-) will not affect the pH of a solution. • An anion that is the conjugate base of a weak acid (ex C2H3O2-) will increase pH. • A cation that is the conjugate acid of a weak base will decrease pH. • The cations of strong bases (group 1A and some of 2A) will not affect pH.

14. More Acidic/Basic Salts • Metal ions besides the cations of strong bases will cause a decrease in pH. • When a solution contains warring ions, then the ion with the larger equilibrium constant will have the greater influence on the pH. • The larger the K value, the stronger the acid/base, so that reaction predominates.

15. Percent Ionization • One whole slide for two equations that mean the same thing! YAAAAYYY! • Percent ionization = concentration ionized[H+]equilibrium original concentration [HA]initial

16. Molecular Structure Don’t you dare fall asleep back there. We’re almost done. All we ask is that you sit through the most boring part of the presentation that we decided to save for last. Gosh. • Binary acids contain H and a single other element (ex. HCl). • Oxyacidscontain one or more O-H bonds. The more electronegative the atom bonded to the central atom, the greater the acid strength. • Carboxylic acids contain one or more carboxyl groups (C-O-OH).

17. The End.