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Acids and Bases

Chapter 21. Acids and Bases. Acids. H +. Acid – a compound that produces ions when dissolved in Examples: Vinegar – Lemon juice – Tea – Ant venom – . water. Acetic acid. Citric acid. Tannic acid. Formic acid. Properties of Acids. S our.

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Acids and Bases

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  1. Chapter 21 Acids and Bases

  2. Acids H+ • Acid – a compound that produces ions when dissolved in • Examples: • Vinegar – • Lemon juice – • Tea – • Ant venom – water Acetic acid Citric acid Tannic acid Formic acid

  3. Properties of Acids Sour • taste • Turns litmus paper • Reacts with metals to form gas • solutions of acids are (must be mixed with water!) • Reacts with to form and Corrosive red H2 Aqueous electrolytes bases H2O salt

  4. Properties of Acids • Sugar, corn syrup, modified corn starch, citric acid, tartaric acid, natural and artificial flavors, yellow 5, yellow 6, red 40, blue 1 • What ingredients make these… so sour?

  5. Naming Binary Acids • H + one other element • Begin with • Use the of the element name • Add the suffix • HCl “hydro-” root “-ic” chlor ic acid hydro

  6. Naming Binary Acids Hydrobromic acid • HBr  • HI  • HF  Hydroiodic acid Hydrofluoric acid

  7. Naming Ternary Acids • H + polyatomic ion • Begin with ion without the • Add suffix if there was an • Add suffix if there was an • HNO3 polyatomic ending -ate -ic -ous -ite Nitrate Nitric acid

  8. Naming Ternary Acids Chloric acid • HClO3  • H3PO3  • H2CO3  Phosphorous acid Carbonic acid

  9. Strength of Acids Strong ionize • acids – completely in water (create a lot of ) • 3 binary acids • Ternary acids Strong if # of atoms - # of atoms ≥ H2SO4 HNO3 H+ HBr HCl HI 2 O H 3 – 1 = 2 4 – 2 = 2

  10. Strength of Acids Weak slightly • acids – ionize only in solution • Binary acids – all others not listed above • Ternary acids Weak if # of atoms - # of atoms ≥ H3PO3 HNO2 aqueous 1 O H 2 – 1 = 1 3 – 3 = 0

  11. Bases OH- • Base – a compound that produces ions when dissolved in • Examples: • Milk of Magnesia – neutralizes stomach acid • Drain cleaner– (hydroxide) water Magnesium hydroxide Sodium hydroxide

  12. Properties of Bases Bitter • taste • Turns litmus paper • solutions of bases are (must be mixed with water!) • Reacts with to form and Slippery blue Aqueous electrolytes acids H2O salt

  13. Naming Bases polyatomic • Use the same rules as for ions (name the cation, then name the anion) • NaOH • Ca(OH)2 • KOH  Sodium hydroxide Calcium hydroxide Potassium hydroxide

  14. Strength of Bases Strong ionize • bases – completely in water (create a lot of ions). • All hydroxides with groups and metals (except ). • bases - only • All bases not listed above as strong. OH- 1 2 Be ionize Weak slightly

  15. Arrhenius Theory Arrhenius H • An acid must contain a and ionize in water to produce • An base must contain a and dissociates in water to produce H+ / H3O+ HCl + H2O  H3O+ + Cl- Arrhenius OH OH- NaOH Na+ + OH-

  16. Disadvantages of Arrhenius Theory OH- • Only compounds with can be classified as a base. What about ammonia, ? • Can only be applied to reactions that occur in • Would classify some compounds as acids, such as NH3 water incorrectly CH4

  17. Arrhenius Acids and Bases • Classify each of the following as an Arrhenius acid (A – acid) or base (A – base). • Ca(OH)2 • HBr  • H2SO4  • LiOH  A - base A - acid A - acid A - base

  18. Bronsted-Lowry Theory acid • A Bronsted – Lowry is any substance that can a • A Bronsted – Lowry is any substance that can a • HCl + H2O  H3O+ + Cl- donate H+ base accept H+ B-L base conjugate base B-L acid conjugate acid Accepts H+ Donates H+

  19. Bronsted-Lowry Theory • Let’s look at the reverse reaction. • Cl- + H3O+ H2O + HCl B-L base conjugate acid B-L acid Donates H+ conjugate base Accepts H+

  20. Bronsted-Lowry Theory Conjugate • acid – formed when a accepts a H+ from an acid. • base – a that remains after an acid gives up a H+. • Conjugate acid – base pair – 2 substances related to each other by the of a single H+. base Conjugate particle accepting/ donating

  21. Types of Acids • Defined by how many H+ they can donate.

  22. Bronsted-Lowry Theory • Identify the acid, base, conjugate acid, and conjugate base. conjugate acid B-L base B-L acid HNO3 + H2O  H3O+ + NO3- Donates H+ conjugate base Accepts H+

  23. Bronsted-Lowry Theory • Give the formula and name of the conjugate base of the following B-L acids. (After the B-L acid donates a H+) • HI  • HCO3- I- Iodide ion Since we take away a +, make the ion more – CO32- carbonate ion

  24. Bronsted-Lowry Theory • Give the formula and name of the conjugate acids of the following B-L bases. (After the B-L base accepts a H+) • H2PO4-  • ClO3- H3PO4 phosphoric acid Since we add a +, make the ion more + HClO3chloric acid

  25. Acidity/Basicity • Water can sometimes act as a B-L acid and sometimes as a B-L base. • The of water: • H2O + H2O  H3O+ + OH- self-ionization conjugate acid B-L base B-L acid conjugate base

  26. Acidity/Basicity • This reaction occurs to a very small extent: = = 1 x 10-7 M [ ] means [H+] x [OH-] = relationship [H+] [OH-] concentration 1 x 10-14 Inverse

  27. Acidity/Basicity pH = 0 pH = 7 pH = 14 [H+] [OH-] acid base neutral

  28. Acidity/Basicity

  29. pH Scale • pHis a measure of how acidic or basic a solution is. • The pH scale ranges from 0 to 14. • Acidic solutions have pH values below 7 • A solution with a pH of 0 is very acidic. • A solution with a pH of 7 is neutral. • Pure water has a pH of 7. • Basic solutions have pH values above 7.

  30. pH Scale • A change of 1 pH unit represents a tenfold change in the acidity of the solution. • For example, if one solution has a pH of 1 and a second solution has a pH of 2, the first solution is not twice as acidic as the second—it is ten times more acidic.

  31. pH • [H+] are often small, so the pH scale is easier to use to represent acidity and basicity. • pH range is from to • log 102 = • log 10-3 = 0 14 pH = -log [H+] 2 -3

  32. pH neutral • In water, a solution, = = pH = [H+] [OH-] 1 x 10-7 pH = negative log [H+] So… take the exponent and change the sign! - (-7) = 7

  33. pH

  34. pH • If [H3O+] = 1.0 x 10 –5 M, what is the pH? • Is the solution basic, neutral, or acidic? Same as [H+] pH = - (exponent) = -(-5) = 5 Because < 7

  35. pH • If [H3O+] = 1.0 x 10 –12 M, what is the pH? • Is the solution basic, neutral, or acidic? pH = - (exponent) = -(-12) = 12 Because > 7

  36. pH • Given that a solution has a pH of 2.0, determine the [H3O+]. pH = - (exponent) 2 = - (exponent) 2 = - (-2) [H3O+] = 1 x 10-2

  37. pOH • Similar to pH, there is also pOH. • Because [H+] x [OH-] = pOH = -log [OH-] 1 x 10-14 pH + pOH = 14

  38. pH/pOH 10-6 10-2 10-4 4 6 12 8 10-12 10-10

  39. pH/pOH 10-8 10-10 10-14 12 10 2 0 6 10-6 10-4 10-2 1

  40. pH/pOH • If [OH-] = 1.0 x 10 –10 M, what is the pOH? • What is the pH? • Is the solution basic, neutral, or acidic? pOH = - (exponent) = -(-10) = 10 pH + pOH = 14 pH + 10 = 14 pH = 4

  41. pH/pOH • What is the pH and the pOH for 1.0 x 10 –6 M HF? • pH • pOH pH = - (exponent) = -(-6) = 6 pH + pOH = 14 6 + pOH = 14 pOH = 8

  42. pH/pOH • Given that a solution has a pH of 9.0, determine the [OH -] and the pOH. • pOH • [OH -] pH + pOH = 14 9 + pOH = 14 pOH = 5 pOH = - (exponent) 5 = - (exponent) 5 = - (-5) [OH-] = 1 x 10-5

  43. pH/pOH cont. • Review: • pH = [H+] = • pOH = [OH-] = • pH + pOH = -log[H+] 10-pH -log[OH-] 10-pOH 14

  44. pH/pOH cont. • Example 1: Determine the pH of a 0.01 M HCl solution. • Example 2: Determine the pH of a 0.0010 M NaOH solution. pH = -log[H+] pH = -log[0.01] pH = 2 pOH = -log[OH-] pOH = -log[0.001] pOH = 3 pH = 11

  45. pH/pOH cont. • Example 3: Determine the pH of a 0.150 M KOH solution. • Example 4: Find [H3O+] for a solution that has a pH of 3.0. pOH = -log[OH-] pOH = -log[0.15] pOH = 0.82 pH = 13.18 [H+] = 10-pH [H+] = 10-3 [H+] = 0.001 M

  46. pH/pOH cont. • Example 5: Find [H3O+] for a solution that has a pH of 8.2. • Example 6: Find [H3O+] and pOH for a solution that has a pH of 4.85. [H+] = 10-8.2 [H+] = 10-pH [H+] = 6.31 x 10-9 M pOH = 14 - pH pOH = 9.15 [H+] = 10-4.85 [H+] = 1.41 x 10-5M

  47. pH/pOH cont. • Example 7: Find [OH-] for a solution that has a pH of 11.2. pOH = 14 – 11.2 = 2.8 [OH-] = 10-pOH = 10-2.8 [OH-] = 1.58 x 10-3 M

  48. What happens with you mix an acid with a base? A ____________________________________________ reaction Neutralization Reactions • What happens with you mix an acid with a base? A reaction • HCl + NaOH + • Products are always a ( and ) and • This is called a reaction double replacement H2O NaCl salt nonmetal metal water neutralization

  49. Neutralization Reactions • Write the balanced chemical equation for the neutralization reaction between: nitric acid and potassium hydroxide + + H2O HNO3 KOH  KNO3 Must use criss-cross to make salt Must balance equation This equation is already balanced

  50. Neutralization Reactions • Write the balanced chemical equation for the neutralization reaction between: sulfuric acid and magnesium hydroxide 2 + + Mg(OH)2 H2O H2SO4 MgSO4  Must balance equation

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