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Acids, Bases, and Indicators

Acids, Bases, and Indicators. Chapter 6, Activity 7. Arrhenius Acids and Bases. Svante Arrhenius (1800s) Defined acids as those compounds that produced hydrogen ions (H + ) in water Examples: HCl – hydrochloric acid H 2 CO 3 – carbonic acid H 2 SO 4 – sulfuric acid.

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Acids, Bases, and Indicators

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  1. Acids, Bases, and Indicators Chapter 6, Activity 7

  2. Arrhenius Acids and Bases • Svante Arrhenius (1800s) • Defined acids as those compounds that produced hydrogen ions (H+) in water • Examples: • HCl – hydrochloric acid • H2CO3 – carbonic acid • H2SO4 – sulfuric acid

  3. Arrhenius Acids and Bases • Bases produce OH- ions in water • Examples: NaOH, Mg(OH)

  4. Concept Test 1 • Decide if the compound is an acid or base: (type the word acid or base, type all three answers, separated by commas): • FrOH • HF • H3PO4

  5. Characteristics • ACIDS • Some react with metal to form H2 gas • Zn + 2HCl -> H2 + ZnCl2 • Corrosive • Sour taste • Turns litmus paper red • Produce H+ in water • BASES • Feel slippery, soapy • Bitter taste • Corrosive • Produce OH- in water PH between 0 and 7 PH between 7 and 14

  6. Turn and Talk • Make a T-chart as below and fill in at least three of each, acid and base, that you might find in your home, school, or workplace. Be prepared to justify your answers:

  7. Example Acid Chemical Equation in water HCl(g) ------> H+(aq) + Cl-(aq) The H+ attaches itself to a water molecule: H+(aq) + H2O(l) -----> H3O+(aq) hydrogen water hydronium ion ion Combining the two gives the full equation: HCl(g) + H2O(l) -----> H3O+(aq) + Cl- (aq)

  8. Example Base Chemical Equation in water NaOH(s) ------> Na+(aq) + OH-(aq)

  9. There are other definitions of acids and bases • Not every acid produces H+ in water • Example: FeCl3 • Not every base produces OH- in water • Example: NH3 • Other definitions are needed • Brønsted-Lowry definition • Lewis definition • More on these later

  10. Neutralization Reactions HCl(aq) + NaOH(aq) -----> NaCl(aq) + H2O (l) • A special type of double-replacement reaction • Acid + base yields a salt plus water

  11. Concept Test 2 • Which of the following is a neutralization reaction? A) KCl + NaNO3 -> KNO3 + NaCl B) HNO3 + KOH -> H2O + KNO3 C) H2O + SO3 -> H2SO4 D) 4Na + O2 -> 2Na2O E) 2NO2 -> 2NO + O2

  12. Concept Test 3 • The neutralization reaction between Al(OH)3 and HNO3 produces the salt with the formula • H2O • AlNO3 • AlH2 • Al(NO3)3 • NO3OH

  13. Concentration • Recall: • Concentration is a measure of how much solute per unit of solvent there is. • Measured in units of molarity (M) • Example: 1.0 MHCl is more concentrated than 0.5 MHCl • REMEMBER! “Strong” does not mean concentrated in chemistry!

  14. What makes and acid or base “strong”? • Disassociates completely in water • Examples: • HCl • HBr • HI • HClO4 • Disassociates completely in water • Examples: • NaOH • KOH • Ca(OH)2 • Sr(OH)2 Strong acids Strong bases

  15. Concept Test 4 Ammonium hydroxide is a weak base because • A) it is a dilute solution. • B) it is only slightly soluble in water. • C) it cannot hold on to its hydroxide ions. • D) it dissociates only slightly in water. • E) it is completely ionized in aqueous solution.

  16. Concept Test 5 Which of the following is correctly identified? • A) NH3, strong acid • B) NaOH, strong base • C) HCl, weak acid • D) H2CO3, strong acid • E) Ca(OH)2, weak base

  17. Titration • You can calculate the concentration of an acid or base solution with a procedure called titration. • EXAMPLE: Find the concentration of an unknown acid solution: • Perform a neutralization reaction with a known concentration of base and an indicator. • When the endpoint is reached, the solution is neutral (the indicator changes color)

  18. What does titration look like?

  19. What does endpoint look like?

  20. Contrast: • Endpoint at pH = 7 • Example: HCl titrated with NaOH • Endpoint at pH higher than 7 • Example: • HC2H3O2 titrated with KOH Strong acid titrated with strong base Weak acid titrated with strong base

  21. Buffers • Solutions that resist changes in pH when a small amount of acid or base is added • Human blood is a natural buffer. Why? (Hint: Your blood should maintain a natural pH of 7.4)

  22. pH Scale

  23. pH Defined • pH = -log[H+] • Thus pOH = -log[OH-] • The brackets indicate concentration in chemistry. • This is a logarithmic scale; for each unit increase, the concentration increases tenfold.

  24. Practice calculation: • What is the pH of a 0.0115 M HCl solution?

  25. Concept Test 6 • What is the pH of a 0.089 M HBrsolution? • -0.05 • Undefined • -0.009 • 1.05 • 13.5 Note that HBr is a strong acid.

  26. Relationship between pH and pOH pH + pOH = 14

  27. Concept Test 7 • What is the pOH of a 0.0815 MNaOH solution? • (Numeric answer) • Note that NaOH is a strong base.

  28. Reminder: • Chem to Go is due tomorrow Chapter 6, Activity 7 Quiz on Friday If you will not be here, plan to take the quiz on Thursday

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