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Ch 13 Acids & Bases

Ch 13 Acids & Bases. Properties Acid-Base Theories Acid-Base Reactions Most assignments are from the Ch 13 handout. Properties. Both conduct electricity (electrolytes) because they break apart to some degree in water. Acids produce H + (proton) in water.

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Ch 13 Acids & Bases

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  1. Ch 13 Acids & Bases Properties Acid-Base Theories Acid-Base Reactions Most assignments are from the Ch 13 handout.

  2. Properties Both conduct electricity (electrolytes) because they break apart to some degree in water. Acids produce H+ (proton) in water. Bases produce OH- (hydroxide) in water. Samples: Acids: vinegar(acetic acid), lactic acid in sour milk, citric acid, Bases: ammonia, lye (NaOH), Milk of Magnesia Mg(OH)2.

  3. More on Acids • Sour taste. NEVER taste acids in lab situations. • Change color of indicators. • Some acids react with metals & release H2 gas. • Acids react with bases to produce salt & water. When neutralization occurs, #1.-#3 disappear. • Conduct electric current.

  4. Acid Nomenclature • Binary Acids contain Hydrogen and another element: • Hydro + root of 2nd element + ic • HF hydrofluoric acid • HCl hydrochloric acid • HBrhdrobromic acid • HI hydroiodic … • H2S hydrosulfuric …

  5. Oxyacids • Contain H, O, and a 3rd element. More are listed in your book.

  6. Common Industrial Acids • Sulfuric • Nitric • Phosphoric • Hydrochloric • Acetic

  7. Bases • Bitter taste (NEVER taste bases in labs). • Change the color of indicators. • Slippery feel (dilute bases, don’t touch concentrated bases) • React with acids to produce salt & water • Conduct electric current.

  8. Neutralization Reaction Acid + Base --> Salt + Water HCl + NaOH --> NaCl + H20 H2SO4 + Ca(OH)2 --> CaSO4 + 2H20 Assignment from Ch 13 A-B handout: 117/1-3 AND 123/1 (naming/formulas)

  9. Arrhenius Acids & Bases • Arrhenius Acid is a chemical compound that increases the concentration of hydrogen ions, H+, in aqueous solutions. • Arrhenius Base is a chemical compound that increases the concentration of hydroxide ions, OH-, in aqueous solutions. • When put with water, these compounds dissociate (break apart) forming ions

  10. HNO3 (l) + H20 (l) --> NO3- (aq) + H30+ (aq) • When put in water, HNO3 , ionizes and the charged particles formed can conduct electricity. • The amount of H30+ (hydronium) produced is an indication of the acid’s strength.

  11. Strong Acids ionize completely in water. Weak Acide release few hydrogen ions in water. Weak Acids: HSO4- H3PO4 HF CH3COOH H2CO3 H2S HCN HCO3- Strong Acids: • HI • HClO4 • HBr • HCl • H2SO4 • HClO3

  12. For Bases, the strength depends on how it dissociates (ionizes) Strong Bases ionize completely. Weal Bases ionize slightly. Weak Bases NH3 + H2O NH4+ + OH- C6H5NH2 “ “ means the reaction is reversible Strong Bases • Ca(OH)2 --> Ca2+ + 2OH- • Sr(OH)2 • Ba(OH)2 • NaOH • KOH • RbOH • CsOH

  13. Assignment • Book: 437/33,34 (strengths & ionization) Answer these on page 118(bottom) of handout. • H/out: 117/5,6 (reactions) and 122/5,6ac (neutralization reactions, balancing and mole ratio)

  14. Acid-Base Theories Bronsted-Lowry Acids donate protons (H+) Molecules or ions can donate protons. HCl + NH3 NH4+ + Cl- Na+ - Na+

  15. The HCl is a Bronsted-Lowry Acid. It donates a proton to water Water can act as a Bronsted-Lowry Acid also as in the following reaction: H2O (l) + NH3 OH- + NH4+

  16. Bronsted-Lowry Bases accept protons. In the equation below, ammonia is the base, because it accepts the proton to become an ammonium ion. acid base HCl + NH3 NH4+ + Cl-

  17. Mono- and Polyprotic Acids • Monoprotic acids can only donate one proton per molecule. Ex.: HCl, HNO3 • Polyprotic acids can donate 2 or more protons per molecule. Ex.: H2SO4, H3PO4 • For polyprotic acids the donations occur in stages, losing one H+ at a time.

  18. Lewis Acids and Bases • Arrhenius and Bronsted-Lowery definitions have some limitations. Lewis classification is based on bonding and structure including substances without hydrogen. The Lewis classification is more complete than the other 2 methods.

  19. A Lewis acid is an atom, ion or molecule that accepts an electron pair to form a covalent bond. Dot notation Structural formula – a bar represents what? A pair of shared electrons.

  20. A Lewis base is an atom, ion, or molecule that donates an electron pair to form a covalent.

  21. Lewis Acid-Base Reaction • is the formation of one or more covalent bonds between an electron-pair donor and an electron-pair acceptor. Pair of donated electrons

  22. Assignment: • 119/1-3 (ionization in stages) • 120/4,5ab (BL a&b, L’LP)

  23. Sample: Dilute HCl(aq) and KOH(aq) are mixed in chemically equivalent quantities. • Write the formula equation for the reaction. HCl(aq) + KOH(aq) --> KCl(aq) + H2O(l) • Write the overall ionic equation. H3O+(aq) + Cl-(aq) + K+(aq) + OH-(aq) --> K+(aq) + Cl-(aq) + 2H20(l) c) Write the net ionic equation. H3O+(aq) + OH-(aq) --> 2H20(l)

  24. Sample: Write the formula equation and net ionic equation for this reaction. Formula equation for: Zn(s) + HCl(aq) --> Zn(s) + 2HCl(aq) --> ZnCl2(aq) + H2(g) Overall ionic equation: Zn(s) + 2H3O+(aq) + 2Cl-(aq) --> Zn2+(aq) + 2Cl-(aq) + H2(g) + 2H20(l) Net ionic equation: Zn(s) + 2H30+(aq) --> Zn2+(aq) + H2(g) + 2H20(l)

  25. Acid-Base Reactions • Now we are going to use Bronsted-Lowry description to explore acid-base reactions. • What was the Bronsted-Lowery theory? • B-L acid donates protons • B-L base accepts protons • Proton: H+ (a hydrogen nucleus)

  26. A conjugate base is the species that remains after a Bronsted-Lowery acid has given up a proton. • A conjugate acid is the species that forms when a Bronsted-Lowery base gains a proton.

  27. Acid-Base Reactions Using Bronsted-Lowry definitions to study Acid-Base reactions, continued: The species that remains after a Bronsted-Lowry acid has given up a proton is the conjugate base of that acid. HF + H2O F- + H30+ Acid conjugate base

  28. The species that is formed when a Bronsted-Lowry base gains a proton is the conjugate acid of that base. HF(aq) + H2O(l) F-(aq) + H30+(aq) Base conjugate acid

  29. HF(aq) + H2O(l) F-(aq) + H30+(aq) Acid Base conjugate conjugate base acid acid1 base2 base1 acid2 Conjugate pairs: • HF and F- • H20 and H30+

  30. Assignment • 121/1,2 (conjugate A&B) • 123/2,3 (same as above)

  31. Strength of Conjugate Acids & Bases • On Page 1 of your handout for this chapter, you have a table which lists and compares the strengths of various acids and their conjugate bases. Get your Ch. 14 handout out now.

  32. Determining direction of equilibrium in Acid-Base reactions The stronger an acid is, the weaker its conjugate base will be. The stronger a base is, the weaker its conjugate acid will be. From these concepts, we can predict the outcome of a reaction.

  33. ) Sample problem on next page:

  34. Sample: Identify the proton donor or acid and the proton acceptor or base. Label each acid-base conjugate pair. CH3COOH + H20 H30+ + CH3COO- acid base conjugate conjugate acid base

  35. Another sample. Write the formula equation, the overall ionic equation, and the net ionic equation for a neutralization reaction that would form RbClO4. Formula equation: RbOH(aq) + HClO4(aq) --> RbClO4(aq) + H20(l)

  36. Overall Ionic equation: Rb+(aq) + OH-(aq) + H30+(aq) + ClO4-(aq) --> Rb+(aq) + ClO4-(aq) + 2H20(l) Net ionic equation: H30+(aq) + OH-(aq) --> 2H20(l)

  37. Amphoteric Compounds These can act as either an acid or a base. Water acts as a base in this reaction: H2SO4(aq) + H20(l) --> H30+(aq) + HSO4-(aq) acid1 base2 acid2 base1 But, water acts as an acid here: NH3(g) + H20(l) NH4+(aq) + OH-(aq) Base1 acid2 acid1 base2

  38. Assignment: • 123/5 (amphoteric/stages) • 124/6ac (conduction, strength)

  39. Review Acids Bases Arrhenius concentration of: [H+] [OH-] Bronsted-Lowry H+ donor H+ acceptor Lewis, e- pair: acceptor donor

  40. Chapter 14, cont.pH and Acid-Base Titration

  41. pH – What is it? pH is an indication of the hydronium ion concentration present in a solution. [H30+] is the symbol for concentration of hydronium ion in moles per liter, mol/L or M pOH is an indication of the hydroxide ion concentration present in a solution. [OH-] is the symbol for concentration of hydroxide ion in mol/L or M

  42. Water self ionizes H20(l) + H20(l) H30+(aq) + OH-(aq) In the above reaction, two water molecules produce a hydronium ion and a hydroxide ion by transfer of a proton. Water is self Ionizing. At 25oC, the concentrations of H30+ and OH- are each only 1.0x10-7 mol/L of water.

  43. Math product of these ions is a constant, kw, the ionization constant of water. Kw= [H30+ ] [OH-] = 1.0x10-7(1.0x10-7) =1.0x10-14 This occurs at 25oC. If the temperature changes, the ion product, Kw changes. When both[H30+ ] and[OH-] are 1.0x10-7, the solution is neutral. If [H30+ ] is greater than 1.0x10-7, the solution is Acidic. (10-6 or 10-4 would be greater) If [OH-] is greater than 1.0x10-7, the solution is Basic.

  44. Calculating without a calculator Kw= [H30+ ] [OH-] = 1.0x10-7(1.0x10-7) =1.0x10-14 Let’s say that the [H30+ ] is 1.0x10-6 and you are asked to find the [OH-]. Kw= [H30+ ] [OH-] --> [OH-] = kw = 1.0x10-14 [H30+ ] 1.0x10-6 -14 – (-6) = -14 + 6 = -8 so: [OH-] = 10-8 mol/Liter More practice: 10-14/10-2 = 10-12 and 10-14/10-9 = 10-5

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