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Chapter 14 Acids and Bases

Learn about the properties, definitions, and reactions of acids and bases. Explore the concepts of Arrhenius and Brønsted and Lowry definitions, electrolytes, strong and weak acids and bases, pH scale, and buffer solutions.

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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

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