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Chapter 14 Acids and Bases PowerPoint Presentation
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Chapter 14 Acids and Bases

Chapter 14 Acids and Bases

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Chapter 14 Acids and Bases

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  1. Chemistry 100 Chapter 14 Acids and Bases

  2. Acids and Bases Acids: sour Bases: bitter or salty

  3. Acids and Bases Arrhenius definition: (If H2O is involved.) Acid: produces H3O+ CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq) H3O+ (Hydronium ion): H+(aq) + H2O(l) H3O+(aq) Base: produces OH- H2O NaOH(s) Na+(aq) + OH-(aq) NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)

  4. Acids and Bases Brønsted and Lowry definition: (If H2O is not involved.) Acid: donates H+ (proton) a proton donor Base: accepts H+ (proton) a proton acceptor HCl + H2O Cl- + H3O+ acid base Conjugate base Conjugate acid Conjugate acid-base pair Conjugate acid-base pair

  5. Acids and Bases HCl + H2O Cl- + H3O+ Proton (H+) is transferred.

  6. CH3COOH + NH3 CH3COO- + NH4+ acid base Conjugate base Conjugate acid Conjugate acid-base pair Conjugate acid-base pair Acids and Bases C6H5OH + H2O C6H5O- + H3O+ acid base Conjugate base Conjugate acid Conjugate acid-base pair Conjugate acid-base pair

  7. Acids and Bases Weak acid and base:is partially ionized in aqueous solution. produces less H+ and OH- CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq) NH3(aq) + H2O(l) NH4+(aq) + OH-(aq) Strong acid and base:is completely ionized in aqueous solution. produces more H+ and OH- HCl(aq) + H2O(l) Cl-(aq) + H3O+(aq) NaOH(aq) + H2O(l) Na+(aq) + OH-(aq)

  8. electrolyte + - Electrolytes bulb Electrolyte:conducts an electric current. Na+ Cl- Ionization (Dissociation) NaCl → Na+ + Cl- strong electrolytes:molecules dissociate completely into ions (NaCl). weak electrolytes:molecules dissociate partially into ions (CH3COOH). nonelectrolytes:molecules do not dissociate into ions (DI water).

  9. Acids and Bases Strong acid/base  Strong electrolyte Weak acid/base  Weak electrolyte HCl(aq) + H2O(l) Cl-(aq) + H3O+(aq)

  10. Acids and Bases A strong acid contains a weak conjugate base.

  11. Amphiprotic:it can act as either an acid or a base. HCl(aq) + H2O(l) Cl-(aq) + H3O+(aq) base NaOH(aq) + H2O(l) Na+(aq) + OH-(aq) acid Acids and Bases Monoprotic acids HCl Triprotic acids H3PO4 Diprotic acids H2SO4

  12. Acids and Bases Oxyacids: acidic H is attached to an oxygen atom. H2SO4 H3PO4 HNO3 Organic acids: contain carboxyl group (-COOH). They are usually weak. CH3COOH

  13. Anion : -ide ion + Hydro -ic acid Naming binary acids HF F-: flouride ion Hydroflouric acid HCl Cl-: chloride ion Hydrochloric acid H2S S2-: sulfuride ion Hydrosulfuric acid

  14. Naming ternary acids -ite ion -ous acid Anion: -ate ion -ic acid HNO2 NO2-: Nitrite ion Nitrous acid HNO3 NO3-: Nitrate ion Nitric acid H2CO3 CO32-: carbonate ion carbonic acid H2SO3 SO32-: sulfurite ion sulfurous acid

  15. [A-] [H3O+] Ka < 1 Ka = K [H2O] = Acid ionization constant [HA] Ka↑ or pKa ↓ Stronger acid Ionization constant HA + H2O A- + H3O+ [A-] [H3O+] K = not for strong acids Equilibrium constant [HA] [H2O] - Log Ka = pKa

  16. Ionization of water H2O(l)+ H2O(l) ⇌ H3O+ (aq) + OH- (aq) KW = [H3O+] [OH-] = (1×10-7) (1×10-7) [H3O+] [OH-] = 1×10-14 pH + pOH = 14

  17. [H+] and [OH-] [H+] = [OH-] Neutral solution [H+] > [OH-] Acidic solution [H+] < [OH-] Basic solution

  18. pH and pOH pH = - log [H3O+] or -log [H+] pOH = - log [OH-] pH scale: 0 7 14 Base Neutral Acid [H3O+] ↑ and [OH-] ↓ [H3O+] ↓ and [OH-] ↑ [H+] = [OH-]

  19. pH meter and pH indicators

  20. Nature & pH indicators Bigleaf Hydrangea In acidic soil In basic soil (alkaline)

  21. pH of strong acids 0.10 M HCl  pH = ? HCl(aq) + H2O(l) Cl-(aq) + H3O+(aq) 1 mol 1 mol 1 mol 0.10 M 0.10 M 0.10 M [H3O+] = [H+] = 0.10 M pH = -log [H+] pH = -log (0.10) = 1.00

  22. Acid Base Reactions Neutralization: reaction between an acid and a base. Acid + Base Salt + Water KOH(aq) + 2HCl(aq) KCl(aq) + H2O(l) 2NaOH(aq) + H2SO4(aq) Na2SO4(aq) + 2H2O(l) Strong acid reacts with strong base to produce the weaker acid and weaker base. (This is the direction of a reaction)

  23. B A Titration (Neutralization reaction) MB: known VB: known MA: unknown VA: known Equivalence point: [H+] = [OH-] when pH = 7 H2SO4(aq) + 2NaOH(aq)  Na2SO4(aq) + 2H2O(l) Base Acid

  24. Titration (Neutralization reaction) Practice: H2SO4(aq) + 2NaOH(aq)  Na2SO4(aq) + 2H2O(l) Find the concentration of NaOH solution if 143 mL of this solution is completely neutralized by 0.205 L of 0.150 M H2SO4 solution? 0.150 mol H2SO4 2 mol NaOH × 1 mol H2SO4 1L H2SO4 solution 0.205 L H2SO4 solution × = 0.0615 mol NaOH solution n 0.0615 mol NaOH M = M = = 0.430 M V (L) 0.143 L NaOH

  25. Buffers Acid or Base pH stays constant. Buffer A buffer resists changes in pH when limited amounts of acid or base are added.

  26. Buffers Our blood is a buffer solution. Acid Acid pH of blood ≈ 7.4 Base Shock Absorber Base

  27. Buffer Composition Weak Acid + its Conjugate base (in equilibrium) salt of the weak acid CH3COOH + CH3COO-Na+ CH3COOH / CH3COO- Or it can be weak base with it’s conjugate acid.

  28. Buffers pH of blood = between 7.35 and 7.45 Carbonate buffer H2CO3 / HCO3- Phosphate buffer H2PO4- / HPO42- Proteins buffer

  29. How do buffers work? Carbonate buffer H2CO3 / HCO3- If we eat an acidic food: HCO3- + H3O+ → H2CO3 + H2O If we eat a basic food: H2CO3 + OH-→ HCO3- + H2O

  30. HA(aq) A-(aq) + H+(aq) pH of Buffers Weak acid Conjugate base [Conjugate Base] pH = pKa + log [Weak Acid] Henderson-Hasselbalch equation [Weak Acid]: concentration of the weak acid [Conjugate Base]: concentration of its conjugate base pKa of the weak acid