1. Acids & Bases

2. Properties of Acids Sour taste Change color of acid-base indicators (red in pH paper) Some react with active metals to produce hydrogen gas Ba (s) + H 2 SO 4(aq) BaSO 4(s) + H 2(g) Some react with bases to neutralize and form salt and water H 2 SO 4 (aq) + 2NaOH (aq) Na 2 SO 4 (aq) + 2H 2 O (l) Some are electrolytes

3. Examples of Acids Lemons and oranges - citric acid Vinegar - 5% by mass acetic acid Pop and fertilizer - phosphoric acid

4. Properties of Bases Bitter taste Change color of acid-base indicators (blue in pH paper) Dilute aqueous solutions feel slippery Ex. Soap Some react with acids to neutralize and form salt and water Some are electrolytes

5. Examples of Bases Soap - NaOH Household cleaners - NH 3 Antacids - Ca(OH) 2, Mg(OH) 2

6. Arrhenius Acids Acids that increase the concentration of hydronium (H 3 O + ) in aqueous solutions HNO 3(aq) + H 2 O (l) H 3 O + (aq) + NO 3 - (aq) H + + NO 3 - + H 2 O acid

7. Why do acids produce H 3 O + ? H + is extremely attracted to the unshared pair of electrons on the water molecule so it donates itself to this molecule where it becomes covalently bonded. The ion formed is known as the hydronium ion (H 3 O + ) H+H+

8. Arrenius Bases Bases that increase the concentration of hydroxide ions (OH - ) in aqueous solutions NaOH (s) Na + (aq) + OH - (aq) H2OH2O

9. Strength of Acids & Bases Strong acids & bases completely ionize in aqueous solutions H 2 SO 4 + H 2 O H 3 O + + HSO 4 - NaOH Na + + OH - Strong acids & bases are strong electrolytes A list of strong acids & bases can be found in your text

10. Weak acids & bases only partially break down into ions when in aqueous solutions HCN + H 2 O H 3 O + + CN - NH 3 + H 2 O NH 4 + + OH - Weak acids & bases are weak electrolytes A list of weak acids & bases can be found in your text

11. Why can we drink H 2 O? Water self ionizes to form equal concentrations of H 3 O + and OH - H 2 O (l) + H 2 O (l) H 3 O + (aq) + OH - (aq) A substance is considered “neutral” when [H 3 O + ] = [OH - ] [H 3 O + ] concentration = 1.0 x 10 -7 M [OH - ] concentration = 1.0 x 10 -7 M

12. When [H 3 O + ] = [OH - ] If [H 3 O + ] > 1.0 x 10 -7 M, the solution is acidic If [OH - ] > 1.0 x 10 -7 M, the solution is basic To find the concentration of [H 3 O + ] or [OH - ] in acidic or basic solutions, the following equation can be used: 1.0 x 10 -14 M 2 = [H 3 O + ] [OH - ] 1.0 x 10 -14 M 2 = ionization constant for H 2 O (K w )

13. Sample Problem A 1.0 x 10 -4 M solution on HNO 3 has been prepared for laboratory use. a. Calculate the [H 3 O + ] of this solution b. Calculate the [OH - ] of this solution c. Is this solution acidic or basic? Why? d. Substitute H 2 SO 4 as the acid. How would the calculations change?

14. Sample Problem An aqueous 3.8 x 10 -3 M NaOH solution has been prepared for laboratory use. a. Calculate the [H 3 O + ] of this solution b. Calculate the [OH - ] of this solution c. Is this solution acidic or basic? Why? d. Substitute Ca(OH) 2 as the base. How would the calculations change?

15. The pH scale The pH scale measures the power of the hydronium ion [H 3 O + ] in a solution The scale typically goes from 1-14 (although it can extend below or above it under extreme conditions) The following equations can be used to determine the pH or [H 3 O + ] of a solution: pH = -log [H 3 O + ] [H 3 O + ] = antilog (-pH) [H 3 O + ] = 1 x 10 -pH

16. pH > 7 basic pH = 7 neutral pH < 7 acidic

17. The pOH scale The pOH scale measures the power of the hydroxide ion [OH - ] in a solution The scale typically goes from 1-14 (although it can extend below or above it under extreme conditions) The following equations can be used to determine the pOH or [OH - ] of a solution: pOH = -log [OH - ] [OH - ] = antilog (-pOH) [OH - ] = 1 x 10 -pOH

18. pH + pOH = 14

19. Sample Problems Calculate the pH of each of the following. Classify as acidic or basic. a.1.3 x 10 -5 M NaOH b.1.0 x 10 -4 M HCl

20. Sample Problems What is the [H 3 O + ] for each of the following? Classify as acidic or basic. a.pH = 5.8 b. pOH = 8.9

21. Sample Problems What is the [OH - ] for each of the following? Classify as acidic or basic. a.[H 3 O + ] = 9.5 x 10 -10 M b.pOH = 1.3

22. Monoprotic versus Polyprotic Acids Monoprotic acids can only donate 1 proton per molecule HCl (g) + H 2 O (l) H 3 O + (aq) + Cl - (aq) Monoprotic

23. Polyprotic acids can donate more than one proton per molecule H 2 SO 4(aq) + H 2 O (l) H 3 O + (aq) + HSO 4 - (aq) Polyprotic HSO 4 - (aq) + H 2 O (l) H 3 O + (aq) + SO 4 -2 (aq) One additional proton can still be donated

24. Strong Acid-Base Neutralization When equal parts of acid and base are present, neutralization occurs where a salt and water are formed HCl (aq) + NaOH (aq) NaCl (aq) + H 2 O (l)

25. Sample Problems H 2 CO 3 + Sr(OH) 2 HClO 4 + NaOH HBr + Ba(OH) 2 NaHCO 3 + H 2 SO 4

26. Titrations When you have a solution with an unknown concentration, you can find it by reacting it completely with a solution of known concentration This process is known as titrating To perform a titration, an instrument called a buret can be used to precisely measure amounts of solution, drop by drop

28. Titration Termonology Equivalence point - the point at which the known and unknown concentration solutions are present in chemically equivalent amounts moles of acid = moles of base Indicator - a weak acid or base that is added to the solution with the unknown concentration before a titration so that it will change color or “indicate” when in a certain pH range (table 16-6 on pg. 495 in your text will show various indicators and their color ranges)

29. End point - the point during a titration where an indicator changes color The 2 most common indicators we will use in our chemistry class will be: Phenolphthalein - turns very pale pink at a pH of 8-10 Bromothymol blue - turns pale green at a pH of 6.2-7.6 Phenolpthalein is clear at pH<8, pale pink at pH 8-10 and magenta at pH >10 Bromothymol blue

30. Practice Titration for an unknown acid 1. Titrate 5.0 of mL of unknown HCl into a 250 mL erlenmeyer flask - *remember to document the starting amount and ending amount of acid on the buret to prevent error 2. Add 2 drops of indicator (phenolphthalein) to the flask - the color of the solution should be clear 3. Titrate with.5M NaOH, continuously swirling the flask, until the solution turns very pale pink for 30 seconds - *remember to document the starting amount and ending amount of base on the buret 4. Mathematically determine the concentration of the unknown HCl solution by using the following equation:

31. Titration Equation M A V A = M B V B M A = molarity (mol/L) of acid V A = volume in L of acid M B = molarity (mol/L) of base V B = volume in L of base moles A = moles B 5. After calculating the molarity of the unknown acid experimentally, get the theoretical molarity and calculate % error

32. Practice titration for an unknown base 1. Titrate 5.0 of mL of unknown NaOH into a 250 mL erlenmeyer flask - *remember to document the starting amount and ending amount of base on the buret to prevent error 2. Add 2 drops of indicator (phenolphthalein) to the flask - the color of the solution should be magenta 3. Titrate with.5M HCl, continuously swirling the flask, until the solution turns very pale pink for 30 seconds - *remember to document the starting amount and ending amount of acid on the buret 4. Mathematically determine the concentration of the unknown NaOH solution by using M A V A = M B V B 5. After calculating the molarity of the unknown base experimentally, get the theoretical molarity and calculate % error