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ACIDS AND BASES

ACIDS AND BASES. Characteristics of Acids. Taste Sour React with metals Electrolytes (conduct electricity) Affect indicators (red=acid) Neutralize Bases pH between 0 and <7. Characteristics of Bases. Taste Bitter Feel Slippery Electrolytes (conduct electricity) Neutralize Acids

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ACIDS AND BASES

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  1. ACIDS AND BASES

  2. Characteristics of Acids • Taste Sour • React with metals • Electrolytes (conduct electricity) • Affect indicators (red=acid) • Neutralize Bases • pH between 0 and <7

  3. Characteristics of Bases • Taste Bitter • Feel Slippery • Electrolytes (conduct electricity) • Neutralize Acids (Antacids) • Affect indicators (base=blue) • pH between >7 and 14

  4. Examples of Acids • HCl monoprotic • H2SO4diprotic • H3PO4 triprotic • HNO3

  5. Examples of Bases • NaOH • Ca(OH)2 • KOH • Soap, Ammonia, • Baking Soda

  6. There are 2 definitions for acids and bases • Arrhenius Theory • Brønsted-Lowry Theory

  7. Arrehenius Theory ON ACIDS…An acid is a substance that ionizes in water to give hydrogen ions (H+)

  8. Arrehenius Theory ON BASES…A base is a substance that ionizes in water to give hydroxide ions (OH-)

  9. An example of the reaction of Hydrochloric Acid: HCl + H2O  H3O+ + Cl- When the HCl breaks apart the free H+is attracted to the H2O and forms a HYDRONIUM ION

  10. An example of the reaction of Potassium Hydroxide: KOH(aq) +  K+(aq) + OH-(aq) When the KOH breaks apart it forms the HYDROXIDE ION (OH-)

  11. So who is the Arrhenius Acid given: H2SO4 + H2O  H3O+ + HSO4- H2SO4 Why? Because it produces H+

  12. So who is the Arrhenius Base given: NaOH + H2O  Na+ + OH- + H2O NaOH Why? Because it produces OH-

  13. The world was cool with the Arrehenius definition until NH3 (Ammonia) came along. When combined with water it did not make OH- ions…looks like we need a new definition.

  14. Brønsted-Lowry Theory ON ACIDS…Acids are substances that are capable of donating a proton (H+)

  15. Brønsted-Lowry Theory ON BASES…Bases are substances that are capable of accepting a proton (H+)

  16. Example of Brønsted Acids and Bases: NH3 + H2O  NH4+ + OH- Here, H2O acts as a Brønsted acid by donating a proton to NH3 which acts as a Brønsted base.

  17. Example of Brønsted Acids and Bases: NH4+ + OH-NH3 + H2O In this case, NH4+ acts as an acid which donates a proton to OH-. OH- acts as a base.

  18. Conjugate Acid The substance formed when a base accepts a H+ ion

  19. Conjugate Base The substance that remains after an acid has donated a H+ ion

  20. Conjugate Acid and Base NH3 + H2O NH4+ + OH- NH4+ + OH-NH3 + H2O For example, NH4+ is the conjugate acid of NH3 (1st equation), and NH3 is the conjugate base of NH4+ (2nd equation).

  21. What is pH? A measurement of the concentration of H3O+ ions (H+ ions)

  22. pH Scale

  23. What is pH? Kw = [H+][OH-] = 1 x 10-14 pH = -log[H+] pOH = -log[OH-]

  24. pH Scale

  25. Neutralization of Bases using Acids

  26. Neutralization acid base nMV = nMV

  27. How many milliliters of 0.80 M HCl are needed to neutralize 35 mL of 1.5 M NaOH? HCl + NaOH  NaCl + H2O 66 mL

  28. What is the concentration of a solution of H2SO4 if it takes 35 mL to neutralize 55 mL of 1.5 M Mg(OH)2? H2SO4 + Mg(OH)2 MgSO4 + 2 H2O 2.4M

  29. How many milliliters of 1.8 M HCl are needed to neutralize 85 mL of 1.0 M NaOH? HCl + NaOH  NaCl + H2O 47 mL

  30. Interesting fact… Water can act as an acid or a base!!!

  31. Lewis Theory ON ACIDS…Acids are substances that are electron-pair receptors

  32. Lewis Theory ON BASES…Bases are substances capable of donating electron-pairs

  33. Amphoteric A substance that can act as either an acid or a base.

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