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Warm-up 1/9/13

Learn the fundamentals of acids and bases, including properties, strengths, ionization constants, pH calculations, neutralization reactions, and pH indicators in liquid solutions.

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Warm-up 1/9/13

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  1. Warm-up 1/9/13 • What do you know about acids and bases? • Daily Objectives: TSW- • Begin acid/base concepts. • Review for final

  2. Unit 11 Acids and Bases

  3. Acid vs Base • All aqueous solutions contain H+ and OH- ions. • Relative amounts determine whether the solution is acid, base, or neutral • Acid soln – more H+ (Hydronium ion) • Basic soln- more OH- • Neutral- equal amounts of each

  4. Which solution is Acidic? Basic? Neutral??? H+ OH- H+ OH- OH- H+ Acidic Neutral Basic Solution Solution Solution

  5. Arrhenius Model • Acid is a substance that contains hydrogen and ionizes to produce hydrogen ions in aqueous solns • Base is a substance that contains a hydroxide group and dissociates to produce a hydroxide ion in aqueous solns • Shortcomings- NH3 (Exception)

  6. Acids and Bases • Produce OH- ions in water • Have a bitter taste and a slippery feel • Break down fats and oils • Formula ends with OH • Poisonous and corrosive to skin • pH greater than 7 • Produce H+ ions in water • Have a sour taste • Break down metals • Formula starts with H • Poisonous and corrosive to skin • pH less than 7

  7. Acid ConjugateBase Brønsted-Lowry Model • Acid- hydrogen ion donor • Base- hydrogen ion acceptor • HX (aq) + H20 (l) H30+(aq) + X-(aq) Base ConjugateAcid

  8. Conjugates • Conjugate Acid • Species produced when a base accepts a hydrogen ion from an acid • Conjugate Base • Species that results when an acid donates a hydrogen ion to a base • Conjugate acid- base pair • Consists of 2 substances related to each other by donating and accepting of a single H+

  9. Conjugates • HF + H2O  H3O+ + F- (H3O+ Conjugate acid) (F- Conjugate base) • NH3 + H20 NH4++ OH- (NH4+ Conjugate acid) (OH- Conjugate base) • Amphoteric- substances that can act as both acids and bases • Monoprotic- HCl, HF • Polyprotic- H2SO4, H3PO4

  10. Acid Strength • Strong acids – ionize completely • Weak acids- do not ionize completely Ka = HCN + H2O  H3O+ + CN- Ka =

  11. Practice Problems • Write an ionization equation and acid ionization constant expression for Nitrous Acid. • HNO2 HNO2 + H2O  H3O+ + NO2- Ka =

  12. One More Practice Problem • Write an ionization equation and acid ionization constant expression for Chlorous Acid. • HClO2 HClO2 + H2O  H3O+ + ClO2- Ka =

  13. Base Strength • Strong Bases- completely dissociate into metal ions and hydroxide ions • Weak bases- partially dissociate Base ionization constant Kb =

  14. Practice Problems Write ionization equations and base ionization constant expressions for the carbonate ion. CO32- CO32- + H2O  HCO3- + OH- Kb =

  15. One More Practice Problem • Write ionization equations and base ionization constant expressions for the hydrogen sulfite ion. • HSO3- HSO3- + H2O  H2SO3 + OH- Kb =

  16. pH • Measure of H+ ions in soln • pH = -log[H+] • Acidic solutions have a pH below 7 • Basic solutions have a pH above 7 • pH 7 is neutral • Change of 1 pH unit represents a tenfold change. (exponential)

  17. pOH • Measures concentration of OH- ion pOH = - log [OH-] pH + pOH = 14.00

  18. Practice Problems • Calculate the pH and pOH of aqueous solutions having the following ion concentrations. [OH-] = 6.5 x 10-6 pOH = -log[OH-] pH = 14.00 – pOH pOH = -log[6.5 x 10-6] pH = 14.00 – 5.19 pOH = -[log 6.5 + log 10-6] ph = 8.81 pOH = -[0.81 + (-6)] pOH = 5.19

  19. One more, one more time! • Calculate the pH and pOH of aqueous solutions having the following ion concentrations. [H+] = 3.6 x 10-9 pH = -log[H+] pOH=14.00 – pH pH = -log[3.6 x 10-9]pOH=14.00-8.44 pH = -[log 3.6 + log 10-9]pOH=5.56 pH = -[0.56 + (-9)] pH = 8.44

  20. Buffers • A buffer is a mixture of a weak acid and its conjugate base OR, a weak base and its conjugate acid. • This mixture resists changes in pH. • The amount of acid or base a buffer can absorb without significant change in pH is called the buffer capacity.

  21. Neutralization Reactions • When an acid is added to a base, the end products are always salt and water. (neutral) • A salt is defined as the neutral end product of an acid/base reaction. ACID + BASE  SALT + WATER H2S + Ca(OH)2  CaS + H2O • What is wrong with this equation???

  22. Balance the final equation! H2S + Ca(OH)2  CaS + H2O 1 Ca 1 1 S 1 4 H 2 2 O 1 H2S + Ca(OH)2  CaS + 2 H2O

  23. Neutralization Reactions Try another example: Acid + Base  Salt + Water H2SO4 + NaOH  Na2SO4 + H2O 1 Na 2 1 SO4 1 3 H 2 1 O 1 H2SO4 + 2 NaOH  Na2SO4 + 2 H2O

  24. Take it one step further… Sulfurous acid and sodium hydroxide yields sodium sulfite and water. H2SO3 + NaOH  Na2SO3 + H2O 1 Na 2 1 SO3 1 3 H 2 1 O 1 H2SO3 + 2NaOH  Na2SO3 + 2H2O

  25. One Last Step Hydrosulfuric acid and calcium hydroxide yields what??? H2S + Ca(OH)2 • One product will always be water. • H2S + Ca(OH)2 H2O + • The other product will be the + ion of the base bonded with the – ion of the acid. H2S + Ca(OH)2 2H2O + CaS

  26. pH Indicators • A chemical substance that changes color in the presence of an acid and/or a base. • 1) pH paper – Dip the paper, match color to scale on vial to determine numeric pH. • pH<7 = acid, pH>7 = base, pH = 7 neutral • 2) Litmus – Dip one red and one blue paper. • Red stays red, blue turns red  Acid • Blue stays blue, red turns blue  Base • Red stays red, blue stays blue  Neutral

  27. pH Indicators 3) Bromthymol Blue – Add a few drops of bromthymol blue to the substance. If the blue color turns to yellow  Acid If the blue color stays blue  Base 4) Phenolphthalein – Add a few drops of phenolphthalein to the substance. If the clear liquid turns to pink  Base If the clear liquid remains clear  Acid

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