1 / 25

Acid-Base Equilibria

Acid-Base Equilibria. Acids and bases are some of the more commonly encountered chemicals Acids and Bases control composition of blood and cell fluids, affect flavors, involved in digestion Bases used in house hold cleaners (NH 3 -based cleansers) Acid rain is an environmental problem

tablita
Download Presentation

Acid-Base Equilibria

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Acid-Base Equilibria Acids and bases are some of the more commonly encountered chemicals Acids and Bases control composition of blood and cell fluids, affect flavors, involved in digestion Bases used in house hold cleaners (NH3-based cleansers) Acid rain is an environmental problem Acids and bases are involved in reactions that produce polymers, synthetic fibers, dyes.

  2. Arrhenius Acid & Base Acid: produces H+ in aqueous solution Base: produces OH- in aqueous solution HCl(aq)  H+ (aq) + Cl- (aq) NaOH(aq)  Na+ (aq) + OH- (aq) Acid + base neutralization: H+(aq) + OH-(aq)  H2O(l) However, an H+ cannot exist by itself in water

  3. Brønsted-Lowry Acids and Bases Acid: proton donor Base: proton acceptor H+: PROTON since the H+ consists of 1 proton and 0 electron HCN(aq) + NH3(aq)  NH4+(aq) + CN-(aq) acid base The H+ is transferred from HCN to NH3 HCN is said to have an acidic H, a hydrogen that can be donated as a H+

  4. HCl has an “acidic” H+, but by itself cannot act as an acid However HCl(aq): HCl (aq) + H2O (l)  H3O+ (aq) + Cl- (aq) H3O+: hydronium ion

  5. HCN - hydrogen cyanide HCN(aq) + H2O(l)  H3O+(aq) + CN-(aq) Only a fraction of HCN donate their H+ to H2O HCN is a weak acid At equilibrium there is both CN- and un-dissociated HCN

  6. In the Brønsted-Lowry theory: a strong acid is fully deprotonated in solution HCl (aq) + H2O (l)  H3O+ (aq) + Cl- (aq) a weak acid is only partially deprotonated in solution HCN(aq) + H2O(l)  H3O+(aq) + CN-(aq) Typically the solvent is water, but not necessarily. An acid that is strong in water, may be weak in another solvent

  7. Brønsted-Lowry Base A proton acceptor. In most cases the molecule possesses a lone pair of electrons to which a H+ can bond to. Example: Oxide, O2- O2- (aq) + H2O(l)  2 OH- (aq) Strong base since all O2- (aq) forms OH- (aq)

  8. NH3: a Brønsted base. The lone pairs on N in NH3 can bond with a H+. NH3 (aq) + H2O (l)  NH4+ (aq) + OH- (aq) NH3(aq) is a weak base; at equilibrium both undissociated NH3 (aq) and NH4+ (aq) exist. A strong base is completely protonated in solution O2- (aq) + H2O(l)  2 OH- (aq) A weak base is partially protonated in solution NH3 (aq) + H2O (l)  NH4+ (aq) + OH- (aq) Strength depends on solvent

  9. Solvent Leveling Since all strong acids are completely de-protonated in water (behave as though they were solutions of H3O+) strong acids are “leveled” in water To compare acidity of acids that are strong acids in water, need to use a solvent in which the “acidity” of the acids differ Strong bases are leveled in water in the same way as strong bases.

  10. Arrhenius definition restricted to water as a solvent However Brønsted-Lowry theory includes non-aqueous solvents CH3COOH (l) + NH3 (l)  CH3COO- (am) + NH4+(am) am - denotes a species dissolved in ammonia Brønsted-Lowry includes acid/base in the absence of solvent Protons can be transferred in the gas phase: HCl(g) + NH3(g)  NH4Cl(s) Acid-base reaction does not have to involve the solvent HCN(aq) + NH3(aq)  NH4+(aq) + CN-(aq)

  11. Conjugate Acids & Bases HCN(aq) + H2O(l)  H3O+(aq) + CN-(aq) acid conjugate base CN- (aq) is the conjugate base of HCN Brønsted-Lowry acids form conjugate bases Acid -----------> conjugate base donates H+

  12. Brønsted-Lowry bases form conjugate acids NH3 (aq) is the base; NH4+ (aq) is the conjugate acid NH3 (aq) + H2O (l)  NH4+ (aq) + OH+ (aq) base -------------> conjugate acid accepts H+

  13. An acid is a proton donor and a base is a proton acceptor. The conjugate base of an acid is the base formed when the acid has donated a proton. The conjugate acid of a base is the acid that forms when the base has accepted a proton.

  14. Lewis Acids & Bases A Lewis base donates a lone pair of electrons A Lewis acid accepts a lone pair of electrons Lewis acids/bases are a broader definition than the Brønsted-Lowry definition H+ is an electron pair acceptor; a Lewis acid Soluble metal oxides are strong bases

  15. NH3 + H2O  NH4+ + OH- base acid Reactions between electron deficient and electron-rich molecules BF3(g) + NH3(g)  F3B - NH3 (s) Lewis acid Lewis base

  16. B-N bond is called a coordinate covalent bond; formed by the coordination of an electron-pair donor to an electron pair acceptor

  17. Amphoterism H2O: acts as both an acid and a base - amphoteric H2O(l) + H2O(l)  H3O+ (aq) + OH- (aq) OH- conjugate base of H2O H3O+ conjugate acid of H2O HCO3- is amphoteric HCO3- (aq) + H2O(l)  H3O+ (aq) + CO32- (aq) HCO3- (aq) + H2O(l)  H2CO3(aq) + OH- (aq)

  18. Water is amphiprotic - both an acid and a base When one molecule transfers a proton to another molecule of the same kind - autoprotolysis or autoionization 2 H2O (l)  H3O+(aq) + OH- (aq) An O-H bond is strong; the fraction of protons transferred is very small.

  19. Calculate the equilibrium constant for the autoionization of H2O(l) 2 H2O (l)  H3O+(aq) + OH- (aq) Kw = [H3O+(aq) ] [OH- (aq) ] DGro = DGfo(H3O+(aq)) + DGfo(OH-(aq)) - 2 DGfo(H2O(l)) = + 79.89 kJ/mol DGro = - R T ln Kw Kw = 1.0 x 10-14 at 298 K

  20. Kw = 1.0 x 10-14 at 298 K Kw = [H3O+(aq) ] [OH- (aq) ] [H3O+(aq) ] [OH- (aq) ] = 1.0 x 10-14 Kw is an equilibrium constant; the product of the concentrations of H3O+ and OH- is always equal to Kw. In pure water [H3O+(aq) ] = [OH- (aq) ] = 1.0 x 10-7 M at 298 K

  21. If the concentration of [OH-(aq)] in increased, then [H3O+(aq)] decreases to maintain Kw.

  22. What are the molarities of H3O+ and OH- in 0.0030 M Ba(OH)2 at 25oC? Ba(OH)2 (aq)  Ba2+ (aq) + 2 OH- (aq) Molarity of [OH- (aq)] = 0.0060 M [H3O+ (aq)] = Kw/[OH- (aq)] = 1.7 x 10-12 M

  23. - - - - pH Scale The concentration of H3O+ can vary over many orders of magnitude A log scale allows a compact description of the H3O+ concentration.

  24. pH = - log [H3O+] [H3O+] = 10- pH mol/L For pure water at 25oC pH = - log (1.0 x 10-7) = 7.00 For a change in pH by 1, H3O+ concentration changes by 10 Higher pH, lower H3O+ concentration pH of pure water is 7 pH of an acidic solution is less than 7 pH of a basic solution is greater than 7

More Related