1. 1 Acids and Bases (courtesy of L. Scheffler, Lincoln High School, 2010)

2. 2 Acids React with certain metals to produce hydrogen gas. React with carbonates and bicarbonates to produce carbon dioxide gas Have a bitter taste Feel slippery. Many soaps contain bases. Bases

3. 3 Properties of Acids Produce H + (as H 3 O + ) ions in water (the hydronium ion is a hydrogen ion attached to a water molecule) Taste sour Corrode metals Good Electrolytes React with bases to form a salt and water pH is less than 7 Turns blue litmus paper to red “Blue to Red A-CID”

4. Some Common Acids HC 2 H 3 O 2 acetic acidin vinegar HClhydrochloric acidstomach acid H 3 C 6 H 5 O 7 citric acidfruits H 2 CO 2 carbonic acidsoft drinks H3 2 PO 4 phosphoric acidsoft drinks

5. 5 Properties of Bases Generally produce OH - ions in water Taste bitter, chalky Are electrolytes Feel soapy, slippery React with acids to form salts and water pH greater than 7 Turns red litmus paper to blue “Basic Blue”

6. Some Common Bases NaOHsodium hydroxidelye KOHpotassium hydroxideliquid soap Ba(OH) 2 barium hydroxidestabilizer for plastics Mg(OH) 2 magnesium hydroxide“MOM” Milk of magnesia Al(OH) 3 aluminum hydroxideMaalox (antacid) Al(OH) 3 aluminum hydroxideMaalox (antacid)

7. 7 Arrhenius Definition Arrhenius Acid - Substances in water that increase the concentration of hydrogen ions ((H + or hydronium ions H 3 O + ). Base - Substances in water that increase concentration of hydroxide ions (OH - ). Categorical definition – easy to sort substances into acids and bases Problem – many bases do not actually contain hydroxides

8. Practice Classify as an acid or a base 1. Taste bitter 2. Taste Sour 3. Feels slimy or slippery 4.Turns litmus paper blue 5. Turns litmus paper red 6. Gives off hydrogen gas when it reacts with some metals 8

9. Practice Classify the following as an Arrehnius Acid or Base and identify what is substance produces in water 1. HNO 3 2. KOH 3. Ca(OH) 2 4. H 2 SO 4 9

10. When neutralization occurs, an acid and a base react together to form a salt and water. Write a balanced equation to represent the neutralization of sulfuric acid and calcium hydroxide, then calculate the mass in grams of calcium hydroxide needed to neutralize 250 mL of 0.01 M solution of sulfuric acid. 10

11. 11 Bronsted-Lowry Definition Acid- substance that donates a proton. Base- substance that accepts a proton. HA + B  HB + + A - Ex HCl + H 2 O  H 3 O + + Cl - Acid Base Conj Acid Conj Base A “proton” is really just a hydrogen atom that has lost it’s electron! The classification depends on how the substance behaves in a chemical reaction

12. Example H 2 SO 4 + NH 3  HSO 4 - + NH 4 + H 2 SO 4 goes to HSO 4 - Did it gain or lose a proton? Is it a BL acid or base? NH3 goes to NH4+ Did it gain or lose a proton? Is it a BL acid or base? 12

13. Identify the BL acid and base 1. HC 2 H 3 O 2 + H 2 O  C 2 H 3 O 2 - + H 3 O + 1. HCO 3 - + HCl  H 2 CO 3 + Cl - 13

14. 14 Conjugate Base-The species remaining after an acid has transferred its proton. Conjugate Acid-The species produced after base has accepted a proton. HA & A - -conjugate acid/base pair A - -conjugate base of acid HA B & HB + -conjugate acid/base pair HB + -conjugate acid of base :B Conjugate Acid Base Pairs

15. A Brønsted-Lowry acid is a proton donor A Brønsted-Lowry base is a proton acceptor acid conjugate base base conjugate acid

16. 16 Note: Water can act as acid or base AcidBase Conjugate Acid Conjugate Base HCl+ H 2 O  H 3 O + +Cl - H 2 PO 4 - + H 2 O   H 3 O + + HPO 4 2- NH 4 + + H 2 O   H 3 O + +NH 3 Examples of Bronsted- Lowry Acid Base Systems

17. 17 Lewis Acid-an electron pair acceptor Base-an electron pair donor Lewis Definition

18. Brønsted-Lowry vs. Lewis All B/L bases are Lewis bases BUT, by definition, a B/L base cannot donate its electrons to anything but a proton (H+) While B/L is most useful for our purposes, Lewis allows us to treat a wider variety of reactions (even if no H+ transfer occurs) as A/B reactions

19. 19 Acid Strength Strong Acid-Transfers all of its protons to water; - Completely ionized; - Strong electrolyte; - The conjugate base is weaker and has a negligible tendency to be protonated. Weak Acid-Transfers only a fraction of its protons to water; - Partly ionized; - Weak electrolyte; - The conjugate base is stronger, readily accepting protons from water  As acid strength decreases, base strength increases.  The stronger the acid, the weaker its conjugate base  The weaker the acid, the stronger its conjugate base

20. 20 Acid Dissociation Constants Dissociation constants for some weak acids

21. 21 Base Strength Strong Base-all molecules accept a proton; - completely ionizes; - strong electrolyte; - conjugate acid is very weak, negligible tendency to donate protons. Weak Base-fraction of molecules accept proton; - partly ionized; - weak electrolyte; - the conjugate acid is stronger. It more readily donates protons.  As base strength decreases, acid strength increases.  The stronger the base, the weaker its conjugate acid.  The weaker the base the stronger its conjugate acid.

22. 22 Common Strong Acids/Bases Strong Bases Sodium Hydroxide Potassium Hydroxide *Barium Hydroxide *Calcium Hydroxide *While strong bases they are not very soluble Strong Acids Hydrochloric Acid Nitric Acid Sulfuric Acid Perchloric Acid

23. A/B Behavior & Chemical Structure 1. Binary Acids Hydrogen and another element 2. Polyprotic Acids Have more than 1 Hydrogen to give away 3. Oxyacids have O in compound 4. Carboxylic Acids have –COOH in compound

24. Wait, water can go both ways? amphoteric substances can behave as either an acid or base depending on what they react with. water and anions with protons (H + ) attached are the most common amphoteric substances

25. Autoionization of Water H 2 O + H 2 O OH - + H 3 O + @ 25 o C the concentrations for both [H 3 O + ] and [OH - ] = 1.00 x 10 -7 and [H 3 O + ] [OH - ] = 1.00 x 10 -14 = K w

26. [H 3 O + ] [OH - ] = 1.00 x 10 -14 = K w Since [H 3 O + ] [OH - ] = 1.00 x 10 -14 = K w [H 3 O + ]=[OH - ] when [H 3 O + ]=[OH - ] the solution is neutral [H 3 O + ]>[OH - ] when [H 3 O + ]>[OH - ] the solution is acidic [H 3 O + ]<[OH - ] when [H 3 O + ]<[OH - ] the solution is basic

27. The pH scale is a way of expressing the strength of acids and bases. Instead of using very small numbers, we just use the NEGATIVE power of 10 on the Molarity of the H + (or OH - ) ion. Under 7 = acid 7 = neutral Over 7 = base

28. pH of Common Substances

29. pH calculations – Solving for H+ If the pH of Coke is 3.12, [H + ] = ??? Because pH = - log [H + ] then - pH = log [H + ] Take antilog (10 x ) of both sides and get 10 -pH = [H + ] [H + ] = 10 -3.12 = 7.6 x 10 -4 M *** to find antilog on your calculator, look for “Shift” or “2 nd function” and then the log button

30. Calculating the pH pH = - log [H+] (Remember that the [ ] mean Molarity) Example: If [H + ] = 1 X 10 -10 pH = - log 1 X 10 -10 pH = - (- 10) pH = 10 Example: If [H + ] = 1.8 X 10 -5 pH = - log 1.8 X 10 -5 pH = - (- 4.74) pH = 4.74

31. pH calculations – Solving for H+ A solution has a pH of 8.5. What is the Molarity of hydrogen ions in the solution? pH = - log [H + ] 8.5 = - log [H + ] -8.5 = log [H + ] Antilog -8.5 = antilog (log [H + ]) 10 -8.5 = [H + ] 3 X 10 -9 = [H + ] pH = - log [H + ] 8.5 = - log [H + ] -8.5 = log [H + ] Antilog -8.5 = antilog (log [H + ]) 10 -8.5 = [H + ] 3 X 10 -9 = [H + ]

32. pOH Since acids and bases are opposites, pH and pOH are opposites! pOH does not really exist, but it is useful for changing bases to pH. pOH looks at the perspective of a base pOH = - log [OH - ] Since pH and pOH are on opposite ends pH + pOH = 14

33. 33 The pH Scale pH [H 3 O + ] [OH- ] pOH

34. pH testing There are several ways to test pH Blue litmus paper (red = acid) Red litmus paper (blue = basic) pH paper (multi-colored) pH meter (7 is neutral, 7 base) Universal indicator (multi-colored) Indicators like phenolphthalein Natural indicators like red cabbage, radishes

35. pH indicators Indicators are dyes that can be added that will change color in the presence of an acid or base. Some indicators only work in a specific range of pH Once the drops are added, the sample is ruined Some dyes are natural, like radish skin or red cabbage

36. 36 Indicators

37. 37 pH and acidity The pH values of several common substances are shown at the right. Many common foods are weak acids Some medicines and many household cleaners are bases.

39. 39 Neutralization An acid will neutralize a base, giving a salt and water as products Examples Acid Base Salt water HCl + NaOH  NaCl + H 2 O H 2 SO 4 + 2 NaOH  Na 2 SO 4 + 2 H 2 O H 3 PO 4 + 3 KOH  K 3 PO 4 + 3 H 2 O 2 HCl + Ca(OH) 2  CaCl 2 + 2 H 2 O A salt is an ionic compound that is formed from the positive ion (cation) of the base and the negative ion (anion) of the acid

40. 40 Titration & Titration Curves Titration: the adding of one solution of an known concentration into another solution standard solution: a solution with a known concentration Titration curve: a graph showing pH vs volume of acid or base added The pH shows a sudden change near the equivalence point The Equivalence point (a.k.a. stoichiometric point) is the point at which the moles of OH- are equal to the moles of H 3 O +

42. 42 Buffer Solutions - Characteristics A solution that resists a change in pH. It is pH stable. A weak acid and its conjugate base form an acid buffer. A weak base and its conjugate acid form a base buffer. We can make a buffer of any pH by varying the concentrations of the acid/base and its conjugate.