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Acids And Bases

Acids And Bases. Chemistry Ms. Piela. Key Characteristics of Acids & Bases. The 3 Main Theories of Acids/Bases. Lewis Acids/Bases. This course will mainly deal with BL theory. Bronsted -Lowry Acids/Bases. Arrhenius Acids/Bases. Theories of Acids & Bases.

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Acids And Bases

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  1. Acids And Bases Chemistry Ms. Piela

  2. Key Characteristics of Acids & Bases

  3. The 3 Main Theories of Acids/Bases LewisAcids/Bases This course will mainly deal with BL theory Bronsted-Lowry Acids/Bases Arrhenius Acids/Bases

  4. Theories of Acids & Bases • Arrhenius Theory of Acids & Bases: • Properties of acids are due to the presence of H+ions • Example: HCl H++ Cl- • Properties of bases are due to the presence of OH- ions • Example: NaOH Na+ + OH-

  5. What is an H+? • H+ ions are bare protons • These are so reactive that they do not exist naturally, but will bond with water to form a hydroniumion, or H3O+ ion • Oftentimes H+ and H3O+ are used interchangeably HCl H++ Cl- HCl(g) + H2O(l)H3O+(aq) + Cl-(aq)

  6. Problems with the Arrhenius theory • Only deals with aqueous solutions (solutions in water) • Not all acids and bases contain H+ and OH- ions • Example: NH3 is a base Considered the most incomplete theory of acids and bases

  7. Theories of Acids & Bases • Brønsted-Lowry Theory of Acids & Bases • Acids are substances that donate H+ ions • Acids are proton (H+) donors • Bases are substances that accept H+ ions • Bases are proton (H+) acceptors • Example: HBr + H2O  H3O+ + Br- A B

  8. Brønsted-Lowry Theory • The behavior of NH3 can be understood now: NH3 (aq) + H2O (l) ↔ NH4+(aq) + OH-(aq) • NH3 becomes NH4+, so NH3 is a proton acceptor (or a Brønsted-Lowry base) • H2O becomes OH-, so H2O is a proton donor (or a Brønsted-Lowry acid)

  9. Brønsted-Lowry Theory

  10. Brønsted-Lowry Theory Conjugate Acid-Base Pairs • Definition: An acid and a base that differ only in the presence or absence of H+ • Every acid has a conjugate base. • Every base has a conjugate acid. • These pairs only ever differ by exactly one hydrogen ion

  11. Brønsted-Lowry Theory • Example Problems • Identify the Brønsted-Lowry acid, base, conjugate acid and conjugate base NH3 + H2O  NH4+ + OH- B A CA CB

  12. Brønsted-Lowry Theory • Example HCl(g) + H2O (l) ↔ H3O+(aq) + Cl-(aq) HSO4- + HCO3- ↔ SO4-2 + H2CO3 A B CA CB A B CB CA

  13. Theories of Acids & Bases • Lewis Acids & Bases • Acids are electron acceptors • Bases are electron donors • Amphoteric – substances that can act as both an acid and a base • Examples: H2O, HCO3-

  14. Summary Of Theories

  15. The pH scale • Developed by SørenSørensen in order to determine the acidity of ales • Used in order to simplify the concept of acids and bases for his workers • The pH scale goes from 0 to 14 • The acidity/basicity of the solutions depends on the concentration of H+ (or H3O+)

  16. The pH scale

  17. pH scale • Low pH values means a high concentration of H+ (acidic) • High pH values means a low concentration of H+ (basic)

  18. Calculations of pH • The Self Ionization of Water • In pure water (pH = 7), the concentrations of the ions (H3O+ and OH-) are equal. [H3O+]=[OH-]= 1x10-7 • This is because water will spontaneously dissociate naturally: H2O (l) ↔ H3O+(aq) + OH-(aq) • Writing the equilibrium expression for the self-ionization of water gives:

  19. The Self-ionization of Water • Plugging in the concentrations in pure water, this gives an equilibrium constant of 1x10-14 • This is referred to as the ion product constant of water • The ion product constant of water has its own symbol: Kw • Unlike other equilibrium constants, the Kw will always be the same value

  20. Calculations of H3O+/OH- • Example #1 • What is the H3O+ concentration in a solution with [OH-] = 3.0 x 10-4 M? Kw = [H3O+][OH-] 1 x 10-14 = [H3O+][3.0 x 10-4] ______ ___________________ 3.0 x 10-4 3.0 x 10-4

  21. Calculations of H3O+/OH- • If the hydronium-ion concentration of an aqueous solution is 1.0 x 10-3 M, what is the hydroxide ion concentration in the solution? Kw = [H3O+][OH-] 1 x 10-14 = [1 x 10-3][OH-] ______ ___________________ 1.0 x 10-3 1.0 x 10-3

  22. Calculations of pH • pH can be expressed using the following equation: pH = -log [H3O+] or [H3O+] = 10-pH • Example #1 • What is the pH of a solution with 0.00010 M H3O+? Is this solution an acid or a base? Acid!

  23. Calculating pH of a solution • Example #2 • What is the pH of a solution where the concentration of hydroxide ions is 0.0136 M? Is this an acid or a base? Kw = [H3O+][OH-] pH = -log [H3O+] Base!

  24. Calculating pH of a solution • Practice #1 • Practice #2

  25. Calculating H3O+/OH- from pH • Example #1 • What is the hydronium ion concentration in fruit juice that has a pH of 3.3? [H3O+] = 10-pH

  26. Calculating H3O+/OH- from pH • What are the concentrations of the hydronium and hydroxide ions in a sample of rain that has a pH of 5.05? [H3O+] = 10-pHKw = [H3O+][OH-]

  27. Calculating H3O+/OH- from pH • Practice #1 • Practice #2

  28. Strength of Acids & Bases • When a solution is considered strong, it will completely ionize in a solution • Nitric acid is an example of strong acid: HNO3 (l) + H2O (l)⇋ NO3-(aq) + H3O+(aq) • In a solution of nitric acid, no HNO3 molecules are present! • Strength is NOT equivalent to concentration!

  29. Strength of Acids & Bases • Knowing the strength of an acid is important for calculating pH • If given concentration of strong acid (such as HNO3) assume it is the same as the concentration of hydronium, H3O+, ions • Given concentration of a strong base, assume it has the same concentration as the hydroxide, OH-, ions

  30. Strong Acids & Bases Ionize 100% 1 M 1 M 1 M • Example NaOH Na+ + OH- OH- Na+ Na+ Na+ OH- OH-

  31. Weak Acids & Bases Ionize X% • Example HF H+ + F- 1 M ?M ?M F- HF H+ H+ HF H+ F- F-

  32. Naming Bases • Bases are soluble metal hydroxides • Follow identical naming rules for ionic compounds • Examples • NaOH • Ba(OH)2 • NH3 • NH4+ Sodium hydroxide Barium hydroxide Ammonia Ammonium

  33. Naming Acids • Binary Acids (HX) • If the acid has an anion that ends in “-ide” use the following basic format to name the acid: • “Hydro – root – ic acid” • Example • HCl Hydrochloric acid

  34. Naming Acids • Example • HBr • Practice • HI • H2S Hydrobromic acid Hydroiodic acid Hydrosulfuric acid

  35. Naming Acids • Polyatomic acids (aka oxoacids, HxAyOz) • Name depends on the polyatomic used: • If polyatomic ends in “-ite”, replace with “ous acid” • If polyatomic ends in “-ate”, replace with “ic acid” • Trick: “I ate something icky”

  36. Naming Acids • Examples • HClO4 • HClO2 • Sulfuric acid Perchloric acid Chlorous acid H2SO4

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