Acids & Bases

1. Acids & Bases Chemistry. Ms. Siddall.

2. Standard 5a: Properties e.x. HCl + Mg  H2(g) + MgCl2

3. Summary 1 • Write the balanced equation for the reaction of aluminum with HCl to form hydrogen and aluminum chloride. • A substance turns red litmus blue and does not react with metal. What is it?

4. Standard 5b: Brønsted-Lowry Definition • Acidsdonate hydrogen ions (Hydrogen ion = H+ = proton) e.x. HBr + H2O  H3O+ + Br- • Basesaccept hydrogen ions e.x. NH3 + H2O  NH4+ + OH-

5. Summary 2 • Write the balanced equation for the reaction of HF (hydrofluoric acid) with water.

6. Strong acid ionization weak acid ionization

7. Standard 5c: Dissociation • Acids: • Strong acid = strong electrolyte • Strong electrolyte = 100% dissociation e.x. HNO3 + H2O  H3O+ + NO3- • in solution (H3O+= hydronium ion) • You MUST know these strong acids: HCl, HBr, HI, HNO3, H2SO4 0% 100%

8. Summary 3 • Which of the following are strong acids? H2SO4, HSO3, HClO2, HBr, HNO3, HNO2

9. Standard 5c: continued • Weak acids = weak electrolytes • Weak electrolyte = Partial dissociation (= equilibrium) e.x. HCN + H2O  H3O+ + CN- • In solution • All other acids are weak acids • Examples: acetic acid HC2H3O2, carbonic acid H2CO3,HCN, H3PO4 ~97% ~3%

10. Summary 4 • At equilibrium does a weak acid solution contain mostly products or mostly reactants? • At equilibrium does a strong acid solution contain mostly products or mostly reactants?

11. Standard 5c: continued Conjugate acids & bases: e.x. HF + H2O  H3O+ + F- • acid • Donates hydrogen ion • Conjugate base • Different from acid by a proton • Base • Accepts hydrogen ion • Conjugate acid • Different from base by a proton CB A B CA

12. Summary 5 • Label the acid (A), base (B), conjugate acid (ca), and conjugate base (cb) in the following reaction: HNO3 + H2O  H3O+ + NO3-

13. Conjugate acids & base pairs: e.x. HF + H2O  H3O+ + F- • Acid and Conjugate base pair = HF & F- • A strong acid has a weak conjugate base • A weak acid has a strong conjugate base • Base and conjugate acid pair = H2O & H3O+ • A strong base has a weak conjugate acid • A weak base has a strong conjugate acid

14. Summary 6 HF + H2O  H3O+ + F- • Identify the acid and conjugate base. Label each as ‘weak’ or ‘strong’ • Identify the base and conjugate acid. Label each as ‘weak’ or ‘strong’

15. Standard 5d: pH scale Definition: pH = -log[H3O+] • pH measures the concentration (amount) of H3O+ ions in solution

16. Summary 7 • Write the equation for HCl reacting with water to make a hydronium ion and a chloride ion • Is the concentration of hydronium ions in this solution high or low?

17. pH0 pH1 pH2 pH3 pH4 pH5 pH6 pH7 pH8 pH9 pH10 pH11 pH12 pH13 pH14 acidic neutral basic

18. Summary 8 Provide an example of: • An acidic compound • A basic compound • A neutral compound Add a minimum of 3 acids and 3 bases to your pH scale

19. Self ionization of pure water H2O + H2O  H3O+ + OH- = Equilibrium system • For pure water: [H3O+] = [OH-] = 1x10-7M • pH = -log[H3O+] = 7 • pH 7 = neutral • [H3O+] > [OH-] = pH < 7= acidic solution • [H3O+] < [OH-] = pH > 7 = basic solution

20. Summary 9 Identify the following solutions as acidic, basic, or neutral: • pH 3 • pH 5 • pH 10 • pH 7

21. Acids ‘produce’ H3O+ in solution e.x. HCl + H2O  H3O++ Cl- • pH measurement: [H3O+] = 0.1M = 1x10-1M pH = -log [H3O+] = -log [1x10-1] pH = 1 • High [H3O+] = ‘Low’ pH reading = acidic solution

22. Summary 10 • pH measures the concentration of ___________________ ions in solution. • A solution with pH = 2: • has a high / low concentration of these ions • is acidic / basic

23. Bases ‘decrease’ H3O+ in solution e.x. NaOH + H3O+ Na+ + 2H2O • pH measurement: [H3O+] = 1x10-14M pH = -log [H3O+] = -log[1x10-14] pH = 14 • Low [H3O+] = ‘High’ pH reading = basic solution (Alkaline)

24. Summary 11 • What is the hydronium ion concentration for a solution with pH = 12? • Is the solution acidic or basic?

26. 5f: calculate pH pH = -log[H3O+] Example: [H3O+] = 4.5 x10-5M • [H3O+] > 1x10-5M • pH should be between 4 - 5 • pH = -log[4.5x10-5] = 4.35

27. Summary 13 • Estimate then calculate the pH for a solution with: • [H3O+] = 7 x 10-2 M • [H3O+] = 2.2 x 10-9 M

28. Relationship between [OH-] and [H3O+] pOH = -log[OH-] pOH + pH = 14 Example: • [OH-] = 1x10-2M • pOH = -log[1x10-2] = 2 • pH = 14 - 2 = 12

29. Summary 14 Complete the following table:

30. 5e: acid/base definitions Arrhenius: • acids are hydrogen containing compounds that ionize to yield H+ ions in aqueous solution. e.x. HCl(aq) H+(aq) + Cl-(aq) • Bases ionize to yield OH- ions in aqueous solutions. e.x. NaOH(aq) Na+(aq) + OH-(aq) • Problems: NH3 is not a base according to Arrhenius

31. Summary 15 • Explain why NH3 is not a base according to the Arrhenius definition of a base.

32. BrØnsted-Lowry: • Acid = hydrogen donor • Base = hydrogen acceptor e.x. NH3(aq) +H2O  NH4+ + OH-

33. Lewis: • An acid accepts a pair of electrons (accepts a negative charge) • Acid = proton donor = electron acceptor • A base donates a pair of electrons • Base = proton acceptor = electron donor • e.x. HCl + H2O  H3O+ + Cl-

34. Electron acceptor

35. Electron donor

36. Summary 16 • Explain the difference between the Bronsted-Lowry definition and the Lewis definition of acids and bases

37. 5g: buffers Buffer = a solution whose pH does not change (much) when acid or base are added. • A buffer is a solution made from an acid and its conjugate base

38. Example: human blood must have a pH between 7.35 – 7.45 (or we will die) • There are 2 buffer systems in human blood: H2PO4-/HPO42- & H2CO3/HCO3- B- + H2PO4- HPO42- + HB B- + H2CO3 HCO3- + HB HB = acid (H+ donor) B- = base (H+ acceptor)

40. Summary 17 • Write a general equation showing the buffering ability of HSO4- and its conjugate base. (use B- as the base that reacts with HSO4-)