1. Acids and Bases Three major ways to define acids and bases introduced by Lewis, Brønsted and Arrhenius. They differ in the role of water Arrhenius and Brønsted require water, Lewis does not Brønsted Acid Donates an H + Base Accepts an H + HCl + H 2 O → H 3 O + + Cl - Acid NaOH + H + → Na + + H 2 O Base

2. Arrhenius Acid Produces H 3 O + when added to water Base Produces OH - when added to water Acids and Bases NH 3 + H 2 O → NH 4 + + OH - HCl + H 2 O → H 3 O + + Cl -

3. Acids and Bases Lewis AcidAccepts electrons Donates electronsBase Note: Electrons are not transferred between acids and bases, they are shared. BH 3 + NH 3 → BH 3 NH 3 Acid Base B H H H N H H H : B N H H H H H H

4. Acids and Bases The Lewis definition is the most general Consider a Brønsted AcidIt donates a H + H + leaves electrons behind A-H A : _ H+H+ i.e. A accepts the electrons A is a Lewis Acid All Brønsted acids are Lewis Acids An Arrhenius acid, is a Bronsted Acid, since it produces H 3 O + when dissolved in water as it “donates” H + to H 2 O.

5. A strong acid, just like a strong electrolyte, is an acid which dissociates completely when dissolved in water. Strong Acids and Bases The concentration of H 3 O + is thereby the highest possible, determined exactly by how much acid was added to water Ex) HCl + H 2 O → H 3 O + + Cl - Ex) H 2 SO 4 + H 2 O → H 3 O + + HSO 4 _ Inorganic acids tend to be strong acids (except HF) A strong base, just like a strong electrolyte, is dissociates completely when dissolved in water. The concentration of OH _ is thereby the highest possible, determined exactly by how much acid was added to water. Ex) NaOH → Na + + OH _ The hydroxides of alkali metals are strong bases.

6. Reactivity A reaction between an acid and a base produces water and a salt HCl (aq) + NaOH (aq) → H 2 O (l) + NaCl (aq) H 3 O + (aq) + Cl - (aq) + Na + (aq) + OH _ (aq) → H 2 O (l) + Na + (aq) + Cl - (aq) H 3 O + (aq) + OH _ (aq) → H 2 O (l) AcidBase WaterSalt Ionic Equation Net equation A strong acid will react completely with any base. A strong base will react completely with any acid.

7. Weak Acids and Bases A weak acid, just like a weak electrolyte, does not dissociate completely when dissolved in water The concentration of H 3 O + is not the highest possible, since much remains in the undissociated form. Ex) CH 3 COOH (l) + H 2 O (l) → CH 3 COO - (aq) + H 3 O + (aq) The concentration of H 3 O + is determined from the dissociation constant, similar to Ksp, and the amount of acid added. Organic acid tend to be weak acids A weak base, just like a weak electrolyte, does not dissociate completely when dissolved in water. The concentration of OH - is not the highest possible, since much remains in the undissociated form. The concentration of OH - is determined from the dissociation constant, similar to Ksp, and the amount of acid added. Ex) NH 3 + H 2 O → NH 4 + + OH - Metal Oxides and nitrogen containing organic compounds tend to be weak bases

9. Non-metal oxides react with water to give oxoacids Amphoteric oxides usually do not dissolve with water by themselves, but react with both strong acids and strong bases to give soluble products The oxides are anhydrides Metal oxides react with water to give hydroxide bases Therefore metal oxides are anhydrides of bases Therefore non-metal oxides are anhydrides of acids Metalloid oxides are amphoteric: they react with either strong acids or strong bases

10. The oxides are anhydrides 1213 14 1516 17 Strength of acids and bases is correlated with the positions of the oxides on the PT MetalMetalliodNon-metal

11. pH The acidity (or basicity) of a solution is reported as pH: pH = -log [H 3 O + ] or [H 3 O + ] = 10 -pH p = power of = concentration of H 3 O + in mol./l = molar (M) For pH < 7 solution is acidic For pH > 7 solution is basic Ex) 0.10 M solution of HCl[H 3 O + ] = 0.10 M pH = - log [0.10] = -(-1.00) = 1.00# of sig. figs. Increased from 2 to 3? For logarithmic quantities only the decimal numbers are significant. Therefore a pH = 1.00 has only 2 sig. figs, Note: pH does not have units

12. pOH Basicity of a solution can be reported as pOH: pOH = -log [OH - ] Where [OH-] = conc. of OH in mol/l pH and pOH are related by: pH + pOH = 14 at 25 o C Therefore an acidic solution as pOH > 7, and a basic solution has pOH < 7 Exercise Determine the pH and pOH of: a) 0.275 M HNO 3 solution [H 3 O + ] = 0.275 M pH = - log (0.275) = -(0.561) = 0.561 pOH = 14.000 – pH = 14.000 -0.561 = 13.439 b) 0.0051 M NaOH solution [OH - ] = 0.0051 M pOH = - log (0.0051) = -(-2.29) = 2.29 pH = 14.000 – pOH = 14.000 -2.29 = 11.71

13. Strength of an Acid The quantity we use to measure the strength of an acid is its pK a (corresponds to how easily an acid gives up H + ). Acids with low pK a values are strong while acids with high pK a values are weak: pK a H 2 O (pK a = 14) HCO 3 -1 (pK a = 10.3) CH 3 CO 2 H (pK a = 4.7) H 3 O + (pK a = 0)  This includes all aqueous solutions of HCl, HBr, HI, HNO 3, HClO 4 and H 2 SO 4 H 3 PO 4 (pK a = 2.15) HF (pK a = 3.1) Concentrated HNO 3 (pK a = -1) Concentrated HCl (pK a = -7) Concentrated H 2 SO 4 (pK a = -3) Citric acid (pK a = 3.1) H 2 CO 3 (pK a = 6.4) STRONG ACIDS WEAK ACIDS

14. Aqua Complexes as Acids A hydrated proton has a pK a of 0 defining the line between the strong and weak acids. How about aqua complexes of metals? What factors affect how easily the aqua complex gives up H + ? CationpK a Approximate pH of a 1 M solution Na(OH 2 ) 6 + 14.27 Ag(OH 2 ) 6 + 126.0 Mg(OH 2 ) 6 2+ 11.45.7 Al(OH 2 ) 6 3+ 52.5 Ti(OH 2 ) 6 4+ -40

15. Aqua Complexes as Acids If we plot pK a versus z 2 /r for a variety of aqua complexes, we see that there is a correlation. If we only look at those metals with low electronegativity values (  1.5), we can approximate: If we introduce an empirical “fudge factor”, we get a more accurate – if more complex formula: Z2rZ2r