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Understanding Acids and Bases: pH Calculation and Neutralization

Learn about the pH scale and how to calculate the concentrations of H+ and OH- ions in a solution. Understand the concept of neutralization and how it relates to the pH of a solution. Explore the different definitions of acids and bases.

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Understanding Acids and Bases: pH Calculation and Neutralization

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  1. Unit 13 Acids and Bases

  2. D. Finding the pH of Solutions Self- ionization of water – the simple dissociation of water H2O H+ + OH- Concentration of each ion in pure water: [H+] = 1.0 x 10-7M + [OH-] = 1.0 x 10-7M Ion-product constant for water (Kw),Where Kw = 1.0 x 10-14 Kw = [H+] [OH-] Acid [H+] > [OH-] Base [H+] < [OH-] Neutral [H+] = [OH-]

  3. Calculating [H+] and [OH-] • reverse the pH equation • The pH of a solution is 7.52. Find the [H+] and [OH-] and determine whether it is acidic, basic, or neutral. • basic

  4. Example 1. If the [H+]in a solution is 1.0 x 10-5M, is the solution acidic, basic or neutral? 1.0 x 10-5 M What is the concentration of the [OH-]? Use the ion-product constant for water (Kw): Kw = [H+] [OH-] 1.0 x 10-14 = [1.0 x 10-5] [OH-] 1.0 x 10-14 = [OH-] 1.0 x 10-5 1.0 x 10-(14-5) pH 5 = acidic 1.0 x 10-9 OH-

  5. Examples 2. If the pH is 9, what is the concentration of the hydroxide ion? Kw = [H+] [OH-] 1.0 x 10-14 = [1.0 x 10-9] [OH-] 1.0 x 10-5 = [OH-] 14 = pH + pOH 14 = 9 + pOH 5 = pOH 3. If the pOH is 4, what is the concentration of the hydrogen ion? Kw = [H+] [OH-] 1.0 x 10-14 = [H+] [1.0 x 10-4] 1.0 x 10-10 = [H+] 14 = pH + pOH 14 = pH + 4 10 = pH

  6. Example • A solution has a pH of 4. Calculate the pOH, [H+] and [OH-]. Is it acidic, basic, or neutral? 14 = pH + pOH 14 = 4 + pOH 10 = pOH • acidic

  7. Practice Problems: Classify each solution as acidic, basic or neutral. 1. [H+] = 1.0 x 10-10 2. [H+] = 0.001 3. [OH-] = 1.0 x 10-7 4. [OH-] = 1.0 x 10-4 Basic pH 10 1.0 x 10-3 acid pH 3 Neutral 14 – 4 = 10 base pH 10

  8. Fill in the chart. 1.0 X 10 -8 1.0 X 10 -6 6 2 12 1.0 X 10 -2 1.0 X 10 -4 4 1.0 X 10 -10 11 3 1.0 X 10 -11 9 1.0 X 10 -5 1.0 X 10 -9 13 1 1.0 X 10 -13

  9. pouvoir hydrogène (Fr.) “hydrogen power” E. pH Scale 14 0 7 INCREASING BASICITY INCREASING ACIDITY NEUTRAL pH is the negative logarithm of the hydrogen ion concentration pH = -log[H+]

  10. E. pH Scale pH = -log[H+] pOH = -log[OH-] pH + pOH = 14

  11. E. The pH Scale

  12. E. pH Scale pH of Common Substances

  13. F. Neutralization • Chemical reaction between an acid and a base. • Products are a salt (ionic compound) and water.

  14. F. Neutralization ACID + BASE  SALT + WATER HCl + NaOH  NaCl + H2O strong strong neutral HC2H3O2 + NaOH  NaC2H3O2 + H2O weak strong basic • Salts can be neutral, acidic, or basic. • Neutralization does not mean pH = 7.

  15. standard solution unknown solution G. Titration • Titration • Analytical method in which a standard solution is used to determine the concentration of an unknown solution.

  16. G. Titration • Equivalence point (endpoint) • Point at which equal amounts of H+and OH- have been added. • Determined by… • indicator color change • dramatic change in pH

  17. G. Titration moles H+ = moles OH- MVn = MVn M: Molarity V: volume n: # of H+ ions in the acid or OH- ions in the base

  18. G. Titration • 42.5 mL of 1.3M KOH are required to neutralize 50.0 mL of H2SO4. Find the molarity of H2SO4. H3O+ M = ? V = 50.0 mL n = 2 OH- M = 1.3M V = 42.5 mL n = 1 MV# = MV# M(50.0mL)(2)=(1.3M)(42.5mL)(1) M = 0.55M H2SO4

  19. Naming Acids • Binary acids • Contains 2 different elements: H and another • Always has “hydro-” prefix • Root of other element’s name • Ending “-ic” • Examples: HI, H2S, HBr, HCl

  20. Naming Acids • Ternary Acids - Oxyacids • Contains 3 different elements: H, O, and another • No prefix • Name of polyatomic ion • Ending “–ic” for “-ate” and “–ous” for “-ite” • Examples: HClO3, H3PO4, HNO2

  21. H2SO3 Sulfurous acid HF Hydrofluoric acid H2Se Hydroselenic acid Perchloric acid HClO4 Carbonic acid H2CO3 Hydrobromic acid HBr Practice

  22. Definitions of Acids and Bases • Arrhenius • Most specific/exclusive definition • Created by Svante Arrhenius, Swedish • Acid : compound that creates H+ in an aqueous solution • Base : compound that creates OH- in an aqueous solution • HNO3 H+ + NO3- • NaOH  Na+ + OH-

  23. Definitions of Acids and Bases • Bronsted-Lowry • A bit more general than Arrhenius definition • Most commonly used definition • Created by two scientists around the same time (1923) • Acid: Molecule or ion that is a proton (H+) donor • Base: Molecule or ion that is a proton (H+) acceptor • HCl + H2O  H3O+ + Cl- • NH3 + H2O ↔ NH4+ + OH-

  24. Definitions of Acids and Bases • Lewis • The most general definition • Defined by electrons and bonding instead of H+ • Created by same scientist who electron-dot diagrams are named after • Acid: atom, ion, or molecule that accepts electron pair to form covalent bond • Base: atom, ion or molecule that donates and electron pair to form covalent bond • NH3 + Ag+ [Ag(NH3)2]1+ • BF3 + F- BF4-

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