Acids & Bases

1. What’s wrong in this picture? Acids & Bases Substances thataffect the pH of solutions.

2. Acids: Are corrosive Taste sour React with indicators Neutralize bases Ex. HCl (hydrochloric acid),H2SO4 (sulfuric acid) Bases: Are slippery Taste bitter React with indicators Neutralize acids Ex. NaOH (sodium hydroxide), NH4OH (ammonium hydroxide) Baking soda (NaHCO3) Acids & Bases typically are, or behave as, IONIC compounds. Litmus is a vegetable dye obtained from certain lichens found principally in the Netherlands.

3. Typical with Acids Typical with Bases The difference between the aqueous solution processes of ionization and dissociation.

4. Acids Most are “hydrogen” bonded with an anion Examples: HNO3(nitric acid) HC2H3O2 (acetic acid) HCl (hydrochloric acid) Bases Most are metal hydroxides Examples: NaOH (sodium hydroxide) KOH (potassium hydroxide) Ba(OH)2 (barium hydroxide) NH4OH* (ammonium hydroxide) Nomenclature

5. Acids Arrhenius - acids donate H+ (in soln) Bronsted-Lowery -acids donate H+ (in soln) Bases Arrhenius - bases donate OH- (in soln) Bronsted-Lowery - bases accept H+ (in soln) Coordinate covalent bond Definitions

6. Conjugate Acid-Base Pairs • The transfer of protons illustrates the characteristics of conjugate pairs • HNO2 + H2O <==> H3O+ + NO2- • NO2- is the conjugate base of HNO2 • H3O+ is the conjugate acid of H2O

7. Monoprotic HCl, HNO3 Diprotic H2CO3 Triprotic H3PO4 Acids can be classified according to the number of hydrogen ions (protons) they can transfer per molecule during an acid-base reaction. Protocity

8. Strong “ions” completely dissociate in water ACIDS: HCl, HBr, HI, HClO4, H2SO4, HNO3 BASES: LiOH, NaOH, KOH, Ca(OH)2, Sr(OH)2, Ba(OH)2 Weak “ions” partially dissociate in water All non-strong acids & bases These exist as equilibrium systems in solution, thus, their “weakness” exists within a range defined by their Keq values. Acid-Base Strength(You can dilute an acid or a base but you can’t change its strength)

9. A comparison of the number of acidic species present in strong acid and weak acid solutions of the same concentration.

10. Two reactions (forward & reverse) occur at the same rate HA <==>H+ + A- BOH <=> B+ +OH- Equilibrium expressions are ways to show the mathematical relationships Keq = [Products]n [Reactants]m n & m are the coefficients of each substance Weak A/B equilibrium

11. HA(aq) + H2O(l) <==> H3O+(aq) + A-(aq) Ka = --------------------- B(aq) + H2O(l) <==> BH+(aq) + OH-(aq) Kb = --------------------- Ionization Constants for Acids & Bases

12. Neutralization reactions - a special type of DR rxn • AX + BY --> AY + BX • HCl + KOH --> HOH + KCl • Acid + Base --> Water + Salt • To balance these rxns. Balance the H in the acid with the OH in the base :)! • For a complete reaction, stoichiometric equivalents of the acid and base must be used.

13. Neutralization equations • HCl(aq) + NaOH(aq) --> H2O(l) + NaCl(aq) • H2SO4 + Ba(OH)2 --> • H3PO4 + KOH --> • HNO3 + Al(OH)3 --> The acid-base reaction between sulfuric acid and barium hydroxide produces the insoluble salt barium sulfate.

14. Calculations • A sample of 0.0084 molHCl is dissolved in water to make 1500 mL solution. Calculate the molarity of the HCl solution and the [H3O+].

15. Water molecules can break apart when they collide H2O(l) <==> H+(aq) + OH-(aq) Kw = ---------------- Kw = 1.0 x 10-14 M2 Adding an acid or a base changes the relative amounts of [H+] and [OH-] but not the value of Kw. Self-Ionization of Water (pH is a derivative of this concept)

16. Ionic Concentration • If [H+] = [OH-] the solution is neutral • If [H+] > [OH-] the solution is acidic • If [H+] < [OH-] the solution is basic • [H+] x [OH-] = 1.0 x 10-14M2 The relationship between H3O+ and OH- in aqueous solution is an inverse proportion.

17. Calculations • If the [OH-] = 3.5 x 10-3 M, what is [H+]?

18. pH: a logarithmic scale of a solution’s hydrogen (hydronium) ion concentration (molarity) • This is a way to express the relative acidity/basicity of a solution. • pH = -log[H+] • Therefore, each difference in pH of 1.0 is equivalent to a concentration change by a factor of 10 • High [H+] causes low pH • Low [H+] causes high pH • Therefore, strong acids have lower pH! • pOH = -log[OH-]

19. pH scale • 0 - 14 is the usual range • pH < 7 = acid • pH > 7 = base • pH = 7 = neutral • pH + pOH = 14

20. Calculations • If the [H+] = 3.35 x 10-5 M, what is the pH of the solution? • On your calculator: - log (3.35 x 10-5) =

21. Calculations • If the [OH-] = 2.8 x 10-4M, what is the pH of the solution?

22. pH --> [H+] calculations • What is the [H+] for a solution with a pH = 3.92? • pH = -log[H+] • 3.92 = -log[H+] • -3.92 =log[H+] • 10-3.92= [H+] • [H+] = 1.20 x 10-4M

23. [H+] = 2.82 x 10-8 M [H+] = 3.98 x 10-11 M [H+] = 7.86 x 10-3 M [H+] = 3.16 x 10-10 M pH = 7.55 pH = 10.4 pH = 2.12 pOH = 4.5 Practice: determine the [H+] for the solutions with the following values.

24. Some aqueous salt solutions have the ability to split (hydrolyze) water and form compounds which result in larger [H+] or [OH-] in the solution. Example: Aluminum chloride AlCl3(aq) --> Al+3(aq) + 3Cl-(aq) Cation of WB Anion of WA Aluminum ion will react with OH- in solution: Remember: H2O<==> H+ + OH- Al+3(aq) + H2O(l) <==> Al(OH)3(aq) + 3H+(aq) Chloride ion will NOT react with H+ in solution! Salt Hydrolysis

25. Rules for Determining pHStrength wins! • Strong Acid + Strong Base --> Neutral sol’n • HCl + NaOH --> NaCl + H2O • Strong Acid + Weak Base --> Acidic sol’n • HCl + Al(OH)3 --> AlCl3 + H2O • Weak Acid + Strong Base --> Basic sol’n • H2S + NaOH --> Na2S + H2O • Weak Acid + Weak Base --> depends on the salt • HNO2 +NH4OH --> NH4NO2 + H2O

26. Buffers • Buffers are solutions in which the pH remains relatively constant whensmallamounts of acid or base are added • Two active chemical species: • A substance to react with & remove added base • A substance to react with & remove added acid. • Buffers are solutions of aweak acid and one of its conjugate baseOR aweak base and one of its conjugate base.

27. Carbonic acid and Sodium bicarbonate • H2CO3 <==> H+ + HCO3- Ka = 1.7 x 10-3 • NaHCO3 --> Na+ + HCO3-

28. Buffering Action in Human Blood • H2CO3 <==> H+ + HCO3- • High concentration High concentration • Ratio: 1 : 10 • Add a base [OH-] and the equilibrium position shifts ; pH doesn’t change much • Add an acid [H+] and the equilibrium position shifts ; pH doesn’t change much • Reason: high [ ] of acid and anion can accommodate large shifts of EQ position. • Lots of acid is produced in thebody daily.

29. Buffer Systems

30. Titration • At the completion of a neutralization reaction (equivalence point) the # moles acid = # moles base So, MaVa = MbVb but, keep the reaction stoichiometry in mind. Diagram showing setup for titration procedures.

31. Chemical Titration • This process can be done for any reaction in which a stoichiometric equivalence is reached and can be identified by an indicator • At the equivalence point an indicator will change color permanently.

32. Calculations • How many mL of 0.10M NaOH solution are needed to neutralize 15 mL of 0.20M H3PO4 solution?