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Chapter 16: Acids and Bases

Chapter 16: Acids and Bases. Students will learn… . 3 definitions of Acids and Bases Acid Strength – pH scale Water as acid and base Calculating pH of strong acids and bases Buffered solutions *Review naming acids, Pg 132 & 133. Three Definitions of Acids-Bases. Arrhenius

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Chapter 16: Acids and Bases

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  1. Chapter 16: Acidsand Bases

  2. Students will learn… • 3 definitions of Acids and Bases • Acid Strength – pH scale • Water as acid and base • Calculating pH of strong acids and bases • Buffered solutions *Review naming acids, Pg 132 & 133

  3. Three Definitions of Acids-Bases • Arrhenius • BrØnsted-Lowry • Lewis

  4. 3 Definitions of Acids and Bases *H+ion always combines with water molecule, forming hydroniumion, H3O+

  5. Arrhenius Acids: HCl, HC2H3O2, HCN • Arrhenius Bases: NaOH, Ca(OH)2 • BrØnsted-Lowry Acids and Bases: can only be determined from chemical equations H2O + HCl→ H3O+ + Cl- base acid conjugateconjugate acid base conjugate pair conjugate pair

  6. Conjugate Acid-Base Pair • An acid and a base that only differ by one H+ ion • Always made of one reactant and one product • Often in reversable reactions (Ex) HSO4- + H2O ↔ SO42- + H3O+ acid base base acid HSO4- + H2O ↔ H2SO4 + OH- base acid acid base

  7. Examples 1. Which of the following represent conjugate acid–base pairs? a) HCl, HNO3 b) H3O+, OH– c) H2SO4, SO42– d) HCN, CN–

  8. 2. Write the conjugate base for each of the following. • HClO4 • H3PO4 • CH3NH3+

  9. Strong vs. Weak • Strong: 100% dissociation of H+ or OH– • Strong electrolytes (conduct electricity very well) • Weak: Less than 100% dissociation of H+ or OH– • Weak electrolytes (conduct electricity somewhat)

  10. Common Strong and Weak Acids/Bases • Strong acids: H+7A, H2SO4, HClO4, HNO3 • Weak acids: HCN, HC2H3O2 • Strong bases: 1A+OH, Sr(OH)2, Ba(OH)2 • Weak bases: NH3 (ammonia) in water = NH4OH (ammonium hydroxide)

  11. Relative Strength of Acids and Bases

  12. Example Acetic acid (HC2H3O2) and HCN are both weak acids. Acetic acid is a stronger acid than HCN. Arrange these bases from weakest to strongest: Cl– CN– C2H3O2–

  13. Terms related Acids • Amphoteric: Ions or molecules that can become an acid and a base (Ex) HSO4-+ H2O → SO42- + H3O+ acid base HSO4-+ H2O → H2SO4 + OH- base acid (Ex) H2O + H2O → H3O++ OH- • Oxyacids: acids containing oxygen (Ex) HNO3, H2SO4, HClO

  14. Monoprotic Acids • Have one proton (H+) to donate (Ex) HCl, HNO3, HI, HClO3 • Diprotic Acids • Have more than two protons to donate (Ex) H2SO4 • Triprotic Acids • Have three protons to donate (Ex) H3PO4

  15. Self-ionization of Water • In neutral water, • H2O + H2O ↔ H3O+ + OH– • [H3O+] = [OH–] = 1×10-7 mol/L • [H3O+]· [OH–] = (1×10-7)·(1×10-7)=1×10-14 @ 25 °= Kw Kw = ion-product constant for water *The constant changes to the temperature

  16. Acidic or Basic Solutions • Kw=1×10-14 always maintained

  17. Common Misunderstanding • An acidic solution has only H+ ions • A basic solution has only OH‒ ions

  18. Example (1) In an acidic aqueous solution, which statement below is correct? a) [H+] < 1.0 × 10–7M b) [H+] > 1.0 × 10–7M c) [OH–] > 1.0 × 10–7M d) [H+] < [OH–]

  19. (2) In an aqueous solution in which [OH–] = 2.0 x 10–10M, the [H+] = ________ M, and the solution is ________. a) 2.0 × 10–10M ; basic b) 1.0 × 10–14M ; acidic c) 5.0 × 10–5M ; acidic • 5.0 × 10–5M ; basic

  20. pH and pOH Scale • pH = –log[H+] • A compact way to represent solution acidity • pH decreases as [H+] increases • [H+] = 10-pH • pOH = ‒log [OH‒] • pOHdecreases as [OH‒] increases • [OH‒] = 10-pOH • pH + pOH = 14

  21. pH = 7; neutral • pH > 7; basic • Higher the pH, more basic. • pH < 7; acidic • Lower the pH, more acidic.

  22. Acidity and pH Scale high pH = weak acid high pOH = weak base

  23. Examples (1) Calculate the pH for each of the following solutions. • 1.0 × 10–4M H+ b) 0.040 M OH–

  24. (2) The pH of a solution is 5.85. What is the [H+] for this solution?

  25. (3) Consider an aqueous solution of 2.0 × 10–3MHCl. What is the pH?  *If the solution was 2.0 × 10–3MHNO2, a weak acid, you can’t take the same approach. Why?

  26. Buffered Solutions • resists a change in its pH when either an acid or a base has been added. • Presence of a weak acid and its conjugate base buffers the solution.

  27. Characteristics of Buffered Solutions • The solution contains a weak acid HA and its conjugate base A–. • The buffer resists changes in pH by reacting with any added H+ or OH– so that these ions do not accumulate. • Any added H+ reacts with the base A–. H+(aq) + A–(aq) → HA(aq) 4. Any added OH– reacts with the weak acid HA. OH–(aq) + HA(aq) → H2O(l) + A–(aq)

  28. Example If a solution is buffered with NH3 to which has been added NH4Cl, what reaction will occur if a strong base such as NaOH is added? a) NaOH + NH3→ NaNH4O b) NaOH + NH4+→ Na+ + NH3 + H2O c) NaOH + Cl–→NaCl + OH– d) NaOH + H2O → NaH3O2

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