Lecture 19.1b- Bronsted-Lowry Acids & Bases

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Section 19.1 lecture (part B) for Honors & Prep Chemistry

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Lecture 19.1b- Bronsted-Lowry Acids & Bases

  1. 1. Bellwork- concentration review How many moles of H+ are in 250ml of 3M HCl? How many moles of OH- are produced when 25g of NaOH is dissolved in 250ml of water? What is the molarity of the NaOH solution?
  2. 2. Arrhenius concept Acids make H+ Bases make OH-
  3. 3. The Brønsted-Lowry definition of ACIDS AND BASES Acids donate protons (H+) HCl  H+ + Cl- Bases accept protons (H+) NH3 + H+  NH4+
  4. 4. The Bronsted- Lowry model is more inclusive than the Arrhenius model. NH3 + H+  NH4+ Ammonia is a Bronsted-Lowry base, but does not dissociate to make OH-
  5. 5. 19.1 Why Ammonia is a Base
  6. 6. HA + H2O  H3O+ + A- Acid base conjugate conjugate acid base An acid donates a proton forming its conjugate base. HA  A- A base accepts a proton forming its conjugate acid. NH3  NH4+
  7. 7. HA  A- Acid conjugate base A- is ready to accept a proton, it is a base. NH3 + H+  NH4+ Base conjugate acid NH4+ has a proton to donate. It is an acid.
  8. 8. 19.1 Conjugate Acids and Bases • A conjugate acid is the particle formed when a base gains a hydrogen ion. • A conjugate base is the particle that remains when an acid has donated a hydrogen ion.
  9. 9. •A conjugate acid-base pair consists of two substances related by the loss or gain of a single proton. •A substance that can act as both an acid and a base is said to be amphoteric.
  10. 10. Water is amphoteric. Water can be an acid or a base H2O  H+ + OH- Water can ionize and donate a proton. H2O  H3O+ As a base, water accepts a proton forming the hydronium ion.
  11. 11. A conjugate acid-base pair consists of two substances related to each other by the donating and accepting of a proton Are the following pairs conjugate acid- base pairs? a. H2O H3O+ b. OH- HNO3 c. HC2H3O2 C2H3O2-
  12. 12. 19.1 Brønsted-Lowry Acids and Bases
  13. 13. Identify conjugate acid base pairs HCl + NH3  NH4+ + Cl- HSO4- + OH-  H2O + SO42- NH3 + H2O  NH4+ + OH-
  14. 14. Lewis Acids and Bases Lewis definition an acid accepts a pair of electrons a base donates a pair of electrons.
  15. 15. 19.1 Lewis Acids and Bases • A Lewis acid is a substance that can accept a pair of electrons to form a covalent bond. • A Lewis base is a substance that can donate a pair of electrons to form a covalent bond.
  16. 16. Animation 25 Compare the three important definitions of acids and bases.
  17. 17. 19.1 Lewis Acids and Bases
  18. 18. 19.1 Section Quiz. 1. Which of the following is NOT a characteristic of acids? a. taste sour b. are electrolytes c. feel slippery d. affect the color of indicators
  19. 19. 19.1 Section Quiz. 2. Which compound is most likely to act as an Arrhenius acid? a. H2O b. NH3. c. NaOH. d. H2SO4.
  20. 20. 19.1 Section Quiz. 3. A Lewis acid is any substance that can accept a. a hydronium ion. b. a proton. c. hydrogen. d. a pair of electrons.
  21. 21. pH The pH scale measures the hydrogen ion concentration[H+] of a solution. A pH of 7 is neutral
  22. 22. A pH less than 7 is acidic (litmus red) A pH greater than 7 is basic (litmus blue) The pH scale ranges from below zero (very acidic) to above14 (very basic)
  23. 23. The pH scale is not linear. The pH scale is logarithmic. pH = -log[H+] [H+] = 1.0 x 10-2 pH = 2 very acidic [H+] = 1.0 x 10-3 pH = 3 acidic A solution with pH of 2 contains 10 times as much H+ as a solution with pH of 3.
  24. 24. Acidic = more H+ than OH- Basic = more OH- than H+
  25. 25. From pH 0 to pH 14 the H+ concentration decreases 100,000,000,000,000 times!!
  26. 26. 19.2 Measuring pH An indicator is a valuable tool for measuring pH because it is a different color in acidic solution than when in base.
  27. 27. 19.2 Measuring pH Phenolphthalein changes from colorless to pink at pH 7–9.
  28. 28. 19.2 Measuring pH
  29. 29. 19.2 Measuring pH Universal Indicators
  30. 30. 19.2 Measuring pH – pH Meters
  31. 31. Solu%on Red Blue pH
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