2. Acids & Bases They are extremely useful in house holds, industry, and everyday life. Some smell pretty, some are pretty, some smell bitter and gross. !! Some are essential to life !! May be hazardous if not used correctly!!!

3. ACIDS Acids have a sour taste Acids contain hydrogen, and some react with active metals to liberate hydrogen gas Acids react with carbonates to produce CO 2 Acids change the color of dyes known as “acid-base indicators” Produce hydrogen ions when they react Acids react with bases to produce salts and water Have pH values less than 7 Acids have formulas that begin with H Acids are electrolytes

4. BASES Taste bitter Feel slippery to the skin Change the color of indicators React to produce hydroxide ions Have pH values greater than 7 Excellent cleaner. Reacts w/ fats and oils React with acids to form salts and water Have formulas that generally end in OH Are electrolytes

5. Svante Arrhenius Swedish chemist 1859 – 1927 ACID- is a chemical compd. That increases the concentration of hydrogen ions (H+) & ionizes in aqueous solutions. H + is best represented by H 3 O + = Hydronium ion. HNO 3 + H 2 0  H 3 O + + NO 3 - HCl + H 2 0  H 3 O + + Cl -

6. Arrhenius Bases Base – increases concentration of hydroxide ions. (OH - ). H 2 O NaOH Na + + OH -

7. Bronsted-Lowry Acid & Base Bronsted-Lowry Acid – is a molecule Or ion that is a proton donor (hydrogen ion) Bronsted-Lowry Base – is a molecule or ion that is a proton acceptor (hydrogen ion). HCl + NH 3  NH 4 + + Cl - Which is the acid? Which is the base?

8. Lewis Acids & Bases 1923 – G. N. Lewis – American Chemist Lewis Acid – atom, ion, molecule that accepts an electron pair to form a Covalent bond. Lewis Base – Atom, ion, molecule that donates a pair of electrons to form a Covalent bond.

9. Strong Acids Hydrobromic acid Sulfuric acid Hydroiodic acid Hydrochloric acid Perchloric acid Nitric acid Phosphoric acid All other acids are weak!!

10. Strong Acids Acid strength increases as polarity increases and bond energy decreases. Strong acids ionize or dissociate completely

11. Strong Bases All hydroxides of group IA and most of IIA are strong bases Strong bases ionize or dissociate completely

12. Weak acids and bases dissociate less than 5 %. Dissociation

13. Monoprotic vs. Diprotic vs.Polyprotic Donates 1 proton Donates 2 protons Donates 2 + protons HCl H2SO4 H3PO4 Triprotic

14. Neutralization Reactions Acids + Bases neutralizes each other to yield a salt + water. Acid + Base Salt + H2O HCl + NaOH NaCl + H2O

15. Neutralization Reactions Acid + Base  HCl + NaOH  HF + NH 3  H 2 SO 4 + KOH  Salt + Water H 2 O + NaCl NH 4 + + F - H 2 O + K 2 SO 4

16. pH Scale – Pouvoir Hydrogène pH is defined as a negative of the common logarithm of the [H 3 O + ]. pH scale ranges from 0 to 14 pH<7 Acid pH=7 Neutral pH>7 Base

17. Indicators change color with pH. Types of indicators Universal indicator pink – acid/ purple – base Litmus paper phenolphthalein Indicators

18. Self Ionization of Water Water can collide with itself to form ions H 2 O + H 2 O  H 3 O + + OH - or H 2 O  H + + OH -

19. KwKwKwKw The ion product constant K w is the same for all solutions. K w = [H + ][OH - ] = 1*10 -14 In acidic solutions [H + ] > [OH - ] In basic solutions [H + ] < [OH - ]

20. Practice If [H + ] = 1*10 -3, what is the [OH - ]? [H + ][OH - ] = 1 *10 -14 [OH - ] = 1*10 -11 M

21. Determining pH pH = - log [H 3 O + ] pH = - log [ 1.0 X 10 –3 M] pH = 3.0 or pH = - log [H + ] -log( 1.0 E -3 )

22. pOH = Base pOH = -log [OH-] If pOH is needed and pH is known, subtract pH from 14. Why remember the pH scale goes from 0 to 14? pH + pOH = 14

23. If [OH - ] = 1*10 -4 M, what is the pH? pOH? pOH = -log[OH - ] pOH = -log(1*10 -4 ) pOH = 4 pH + pOH = 14 pH + 4 = 14 pH = 10

24. If pH = 6.0, what is the [H + ]? pOH? pH = -log[H + ] 6.0 = -log[H + ] -6.0 = log[H + ] Antilog(-6.0) = antilog (log[H + ]) 1*10 -6 =[H + ] pH + pOH = 14.0 6.0 + pOH = 14.0 pOH = 8.0

25. If pH = 5, what is the [OH - ]? pH + pOH = 14.0 5.0 + pOH = 14.0 pOH = 9.0pOH = -log[OH - ] 9.0 = -log[OH - ] -9.0 = log[OH - ] Antilog(-9.0) = antilog (log[OH - ]) 1*10 -9 =[OH - ]

26. Practice calculating pH What is the pH of the solution that has 1.0 X 10 –3 M. ? The pH of human blood is 7.41. Calculate the[H 3 0 + ] pH = - log [H 3 0 + ] antilog(-7.41)= 3.89 X 10 -8 A solution. @ 25 C has a hydronium concentration of 1.0 X 10 –9. What is the pH level?

27. One consequence of acid rain is vividly illustrated by these two photographs f a forest site in Germany. The one on the left was taken in 1970, the one on the right 1983. Acid rain in this region has now been reduced, as a result of better control of power plant emissions.

28. Common Acids Sulfuric Acid –Fertilizer, refining, metallurgy, and battery acid –Add to water to liberate heat Phosphoric Acid-flavor Nitric Acid-explosives Hydrochloric Acid- 0.4% in stomach Acetic Acid- 4% to 8% in vinegar

29. Conjugate Acids and Bases

30. Titrations The controlled addition and measurement of the amount of a solution of a known concentration that is required to react completely with a measured amount of a solution of unknown concentration Standard solution -a solution that contains a precisely known concentration of a solute