Acids & Bases

1. Acids & Bases

2. Key Characteristics of Acids & Bases

3. Theories of Acids & Bases • Arrhenius Theory of Acids & Bases • Properties of acids are due to the presence of H+ ions • Example: HCl H++ Cl- • Properties of bases are due to the presence of OH- ions • Example: NaOH Na+ + OH-

4. H+ ions in water • H+ ions are bare protons • These H+ ions react strongly with the nonbonding pair of electrons in a water molecule • This forms the hydroniumion, H3O+ • Oftentimes H+ and H3O+ are used interchangeably HCl H+ + Cl- HCl(g) + H2O(l)H3O+(aq)+ Cl-(aq)

5. Problems with Arrhenius • Arrhenius theory has limitations: • Only deals with aqueous solutions (solutions in water) • Not all acids and bases produce H+ and OH- ions • NH3 for example is a base • Brønsted and Lowry proposed a definition based on acid base reactions transferring H+ ion from one substance to another

6. Brønsted-Lowry Theory

7. Theories of Acids & Bases • Brønsted-Lowry Theory • Acids are substances that donate H+ ions • Acids are proton donors • Bases are substances that accept H+ ions • Bases are proton acceptors • Example: HBr + H2O  H3O+ + Br- AB

8. Brønsted-Lowry Theory • The behavior of NH3 can now be understood: NH3 (aq) + H2O (l) ↔ NH4+(aq) + OH-(aq) • Since NH3 becomes NH4+, it is a proton acceptor (or a Brønsted-Lowry base) • H2O becomes OH-, which means it is a proton donor (or a Brønsted-Lowry acid)

9. Brønsted-Lowry Theory Conjugate Acid-Base Pairs • An acid and a base that differ only in the presence or absence of H+ are called a conjugate acid-base pair. • Every acid has a conjugate base. • Every base has a conjugate acid. • HX is the conjugate acid of X- • H2O is the conjugate base of H3O+

10. Brønsted-Lowry Theory • These pairs differ by only one hydrogen ion • Example • Identify the Brønsted-Lowry acid, base, conjugate acid and conjugate base NH3 + H2O  NH4+ + OH- BACACB • NH3 acts as a Brønsted base by accepting a proton. • Water acts as a Brønsted acid by donating a proton.

11. Brønsted-Lowry Theory • Example HCl(g) + H2O (l) ↔ H3O+(aq)+ Cl-(aq) HSO4- + HCO3- ↔ SO4-2 + H2CO3 A B CA CB A B CB CA

12. Theories of Acids & Bases • Lewis Acids & Bases • Acids are electron acceptors • Bases are electron donors • Example: H2O + NH3 OH- + NH4+ • Is really: H2O + :NH3 OH- + H:NH3+ Electron pair donor(NH3) Electron pair acceptor(H+)

13. Summary Of Theories

14. The Self-Ionization of Water • Even pure water contains a small number of ions: H2O (l) ↔ H3O+(aq) + OH-(aq) • In pure water, the concentrations of the ions (H3O+ and OH-) are equal. [H3O+]=[OH-]= 1x10-7 M

15. The Self-ionization of Water • Writing the equilibrium expression for the self-ionization of water gives: • Plugging in the concentrations in pure water, this gives an equilibrium constant of 1x10-14 • this is referred to as the ion product constant of water • This ion product constant of water is given the symbol Kw

16. The Self-ionization of water • Example #1 • What is the H3O+ concentration in a solution with [OH-] = 3.0 x 10-4 M? Kw = [H3O+][OH-] 1x10-14 = [H3O+][3.0x10-4]

17. Example #2 • If the hydroxide-ion concentration of an aqueous solution is 1.0 x 10-3 M, what is the [H3O+] in the solution? Kw = [H3O+][OH-] 1x10-14 = [H3O+][1.0x10-3]

18. The pH scale • Developed by SørenSørensen in order to determine the acidity of ales • Used in order to simplify the concept of acids and bases • The pH scale goes from 1 to 14 • A change in one pH unit corresponds to a power of ten change in the concentration of hydronium (H3O+) ions • A pH = 2.0 has 10 times the concentration of H3O+ than a pH = 3.0, and 100 times greater than pH = 4

19. The pH scale

20. Calculations of pH • pH can be expressed using the following equation: pH = -log [H3O+] or [H3O+] = 10-pH • Example #1 • What is the pH of a solution with 0.00010 M H3O+? Is this solution an acid or a base? Acid

21. Calculations of pH • Example #2 • What is the pH of a solution with the concentration of hydroxide ions 0.0136 M? Is this an acid or a base? pH = -log [H3O+] Kw = [H3O+][OH-] Base

22. Calculations of pH • Practice #1 • Practice #2

23. Calculations of pH • Example #1 • What is the hydronium ion concentration in fruit juice that has a pH of 3.3? [H3O+] = 10-pH

24. Calculations of pH • What are the concentrations of the hydronium and hydroxide ions in a sample of rain that has a pH of 5.05? [H3O+] = 10-pHKw = [H3O+][OH-]

25. Calculation of pH • Practice #1 • Practice #2

26. Strength of Acids & Bases • When a solution is considered strong, it will completely ionize in a solution • Nitric acid is an example of strong acid: HNO3 (l) + H2O (l) NO3-(aq) + H3O+(aq) • In a solution of nitric acid, no HNO3 molecules are present • Strength is NOT equivalent to concentration!

27. Strength of Acids & Bases • Knowing the strength of an acid is important for calculating pH • If given concentration of strong acid (such as HNO3) assume it is the same as the concentration of hydronium, H3O+, ions • Given concentration of a strong base, assume it has the same concentration as the hydroxide, OH-, ions

28. Strong Acids & Bases Ionize 100% 1 M 1 M 1 M • Example NaOH Na+ + OH- OH- Na+ Na+ Na+ OH- OH-

29. Weak Acids & Bases Ionize X% • Example HF H+ + F- 1 M ?M ?M F- HF H+ H+ HF H+ F- F-

30. Strength of Acids & Bases

31. Strength of Acids & Bases

32. Strong Acids • Must be memorized!

33. Strong Acids • 6 of 7 strong acids are monoprotic (HX) • Exists only as H ions and X ions HI(aq) H+(aq) + I-(aq) 2M HI = [H+]= [I-] = 2M • Determining pH of Strong Acids • For Strong Acids: pH = -log [H+] • For monoprotic strong acids: [H+] = [X]