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Understanding Acids and Bases: Properties, Strengths, and Theories

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kalli
58 diapositives · 747 Views · Jul 21, 2014

This chapter explores the essential concepts of acids and bases in chemistry. It defines acids as compounds that produce H+ ions when dissolved in water, with examples such as vinegar and lemon juice. The properties of acids and bases, their strengths, and the rules for naming binary and ternary acids are outlined. The chapter also discusses the Arrhenius and Bronsted-Lowry theories, highlighting how substances classify as acids or bases based on their ability to donate or accept protons. Additionally, it covers the self-ionization of water and the concept of pH.

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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