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Understanding Acids and Bases: Properties, Definitions, and Electrolytes

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This chapter explores the properties and definitions of acids and bases, outlining key characteristics such as taste, pH levels, and reactions with indicators. It distinguishes between strong and weak acids and bases, discussing their ionization in solutions and their classifications as electrolytes. Furthermore, it introduces concepts such as polyprotic acids and neutralization reactions, explaining how to determine if a solution is acidic, basic, or neutral. A focus on electrolytes highlights their role in conducting electrical current, emphasizing the importance of understanding these concepts in chemistry.

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Understanding Acids and Bases: Properties, Definitions, and Electrolytes

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  1. Chapter 15 Acids and Bases

  2. Objectives • Describe the properties of acids and bases • Recognize acids and bases by definitions • Compare strong and weak acids/bases • Describe electrolytes

  3. What Do You Know About Acids and Bases?

  4. Properties of Acids • Sour taste • (A VERY bad way to test in the lab!) • What acids have you tasted? • Turns blue litmus red • Reacts with metal to form hydrogen gas • pH less than 7 • Electrolyte • Solutions that conduct electrical current

  5. Properties of Bases • Tastes bitter • (Also not a good lab practice!) • Turns red litmus blue • pH greater than 7 • Electrolyte

  6. Acid/Base Definitions • Arrhenius Definition • Acids increase the H+ ion concentration in solution • Acids have H as the first element • Bases increase the OH- ion concentration in solution • Many bases end with OH • Definition limited to solutions!

  7. Acid/Base Definitions • Bronsted/Lowry Definition • Acids are H+ donors • H+ ion is just a proton • (often called proton donors) • Bases are H+ acceptors • (Often called proton acceptors) • This is the definition of acids and bases we will use the most.

  8. Strong Acids • Strong Acids – Acids that completely ionize in solution • Assume that HA can be an acid HA  H+ + A- • There will be NO HA left in solution • Strong Acids – Memorize the first 3 HNO3 HCl H2SO4 HClO4 HBr HI

  9. The Hydronium Ion • H+ ions can be written one of two ways • 1st H+ • Easy way to indicate the ion • 2nd H3O+ • Indicates that the H+ attaches to water molecules • Hydronium Ion • Either way is fine

  10. Weak Acids • Weak Acids – Acids that only partially ionize in solution • Equilibrium is established • Assume HA to be a weak acid HA + H2O  H3O+ + A- • There will be lots of HA left in solution • Weak Acids – Any acid that is not strong H2CO3 HF H3PO4 HNO2 HC2H3O2 H2SO3

  11. Inorganic vs Organic Acids • All acids have H in the formula • Inorganic acids contain hydrogen and a halogen or hydrogen and a polyatomic ion • Organic acids have hydrogen and carbon and oxygen (not CO3-2) • Carboxyl group • Carboxylic acids • All weak

  12. Strong Bases • Strong Bases – Bases that completely ionize in solution KOH  K+ + OH- • There will be NO KOH left in solution • Strong Bases NaOHKOH • Group 2 hydroxides are also considered strong bases

  13. Weak Bases • Weak Bases – Bases that partially ionize in solution NH3 + H2O  NH4+ + OH- • There will be lots of NH3 left in solution • Weak bases are ammonia derivates • They will have nitrogen in them

  14. Electrolytes • Conduct current because charged particles are free to move about • Acids form ions when they are dissolved in solution HCl  H+ + Cl- • Salts form ions when they are dissolved in solution NaNO3  Na+ + NO3- • Charged particles complete (or close) the circuit

  15. Closed Circuit

  16. Open Circuit

  17. Another Type of Open Circuit

  18. Closed circuit • Ions in solution close the circuit • Charges flow to the opposite pole • Current flows from negative to positive • Light bulb shines!

  19. Nonelectrolytes • Solutions that do not conduct an electric current • No ions are present • Nothing to conduct the current • All molecules are nonelectrolytes • They dissolve into individual molecules

  20. Electrolytes or Nonelectrolytes? • CH4 • Nonelectrolyte • KBr • Electrolyte • C2H6O • Nonelectrolyte • H2SO4 • Electrolyte

  21. Strong and Weak Electrolytes • Strong electrolytes • Solutions that conduct current well • Lots of ions in solution • Strong acids and bases, salts • Weak electrolytes • Solutions that conduct current poorly • Few ions • Weak acids and bases

  22. Homework • p. 625 #42,43,51,55,61,67,70,86,97

  23. Objectives • Recognize polyprotic acids. • Compare the strengths of weak acids and bases. • Describe a neutralization reaction • Calculate neutralization data • Determine if a solution is acidic, basic, or neutral • Explain the autodissociation of water

  24. Polyprotic Acids • Acids that can donate more than 1 hydrogen • H2CO3 – Carbonic Acid • Able to donate 2 H’s • Diprotic • H3PO4 – Phosphoric Acid • Able to donate 3 H’s • Triprotic

  25. Cont. • HNO3 – Nitric Acid • Only 1 H to donate • Monoprotic • HC2H3O2 – Acetic Acid • Only the first H can be donated • Same with most carboxylic acids

  26. How Weak is Weak?Quantifying Weak Acids and Bases • Weak acids and bases have equilibrium dissociation values • The smaller the constant the weaker the acid or base • The larger the constant the stronger the acid or base

  27. How Weak is Weak?Quantifying Weak Acids and Bases Acid - HA(aq) H+ + A- • The Equilibrium Expression Ka = [H+] [A-] / [HA] Base - B(aq) + H2O  BH+ + OH- • The Equilibrium Expression Kb = [BH+] [OH-] / [B]

  28. Sample Ka Values • Acid Ka • HClO2 1.2x10-2 • HF 7.2x10-4 • HC2H3O2 1.8x10-5 • HClO 3.5x10-8 • HCN 6.2x10-10 • HIO 2.0x10-11 • What is the strongest / weakest acid?

  29. Sample Kb Values • Base Kb • NH3 1.8x10-5 • (C2H5)2NH 1.3x10-3 • (C2H5)3N 4.0x10-4 • CH3NH2 4.4x10-4 • C2H5NH2 5.6x10-4 • What is the strongest / weakest base?

  30. Neutralization Reactions • The reaction of an acid and a base to yield a salt and water • HI + KOH  HOH + KI • Net Ionic Equation • Net Ionic is ALWAYS the same

  31. Example • What volume of 0.25M KOH is required to react with 50.0 mL of 0.20 M HBr?

  32. Example • What volume of 1.2M NaOH is required to react with 30.0 mL of 0.70 M H2SO4?

  33. Acidic, Basic, and Neutral Solutions • Relates to concentrations of H+ and OH- • Acidic solutions have more H+ than OH- • H+ from an acid • Basic solutions have more OH- than H+ • OH- from base • Neutral solutions have equal amounts of both ions • Both ions come from water!

  34. Autoionization of Water • Water molecules dissociate by themselves H2O(l) H+ + OH- • In pure water the concentrations are equal • Both 1.00x10-7M • Write the equilibrium expression • Kw is always 1.00x10-14 at 25ºC

  35. LeChatlier’s Principle (Again) • The concentration of H+ and OH- can vary in solution. • When H+ is added, [OH-] decreases • When OH- is added, [H+] decreases • That is how solutions become acidic or basic

  36. Example • What is the concentration of H+ and OH- in a 0.10M solution of HCl? Is the solution acidic, basic, or neutral?

  37. Example • A solution is prepared by dissolving 6.00 grams of KOH in 200.0mL of water. What is the concentration of H+ and OH-? Is the solution acidic, basic, or neutral?

  38. Homework • p. 626 #80,87,89,92,101,117,119,124,131,134

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