1. Acids and Bases Chapter 19

2. Ions in Solution  Aqueous solutions contain H + ions and OH - ions  If a solution has more H + ions than OH - ions it is acidic  If a solution has more OH - ions than H + ions it is basic  If a solution has the same number of OH - ions and H + ions it is neutral

3. Properties of Acids  Properties acids –  Produce H + ions when dissolved in water  taste sour  Turn Blue Litmus paper Pink  React with metals to produce Hydrogen Gas  good at dissolving things (food in stomach, teeth to form cavities, mineral deposits in coffeemaker)  Have a pH of 0 to < 7  Conduct electricity

4. Properties of Bases  Properties bases –  taste bitter  feel slippery  tend to produce OH - ions when placed in water  turn Red Litmus Paper Blue  Have a pH of 7 to 14  Conduct electricity

5. Two models of Acids and Bases  Arrhenius  Bronsted-Lowry  Two similar but different explanations, both are correct

6. Arrhenius  Acid - produces H + in aqueous solution  HCl (g)  H + (aq) + Cl - (aq)  Base – produces OH - in aqueous solution  NaOH (s)  Na + (aq) + OH - (aq)  Explains most Acids and Bases, but not all

7. Acid-Base Neutralization Reaction  When an acid and a base are mixed together: acid + base   A salt and water are formed:  a salt + water  remember a salt is any ionic compound

8. Examples  HCl (aq) + NaOH (aq)  NaCl (aq) + H 2 O (l)  Acid+ BaseA Salt + Water  2HCl(aq) + Ca(OH) 2 (aq)  CaCl 2 (aq) + 2 H 2 O(l)  Acid+ BaseA Salt + Water

9. Bronsted-Lowry Model  Ammonia (NH 3 ) is a base.  It does not contain OH - in it.  So according in Arrhenius it isn’t a base.  But when placed in water it produces OH, so it must be a base.

10. Bronsted-Lowry Model  The Bronsted-Lowry model focuses on H + ion  Acid – H + donor  When placed in water Acids give H + away  HCl (g) + H 2 O (l)  Cl - (aq) + H 3 O + (aq)  H 3 O + (aq) is called the Hydronium ion  Base – H + acceptor  NH 3 (aq) + H 2 O (l)  NH 4 + (aq) + OH -

11. Acid-Base Pairs  In Bronsted-Lowry Acids and Bases you must always have a pair  Every Acid must have a Conjugate Base  Every Base must have an Conjugate Acid HCl (g) + H 2 O (l)  Cl - (aq) + H 3 O + (aq) Acid Base Con Base Con Acid  HCl is an Acid and produces Cl -  H 2 O must be a Base and produces H 3 O +

12. More on Pairs  HI(s) + H 2 O(l)  H 3 O + (aq) + I - (aq)  AcidBase  C.A. C.B.  H 2 SO 4 (s) + H 2 O(l)  H 3 O + (aq) + HSO 4 - (aq)  Acid Base  C.A. C.B.

13. Strong Acids  Strong acid - reacts completely with water to produce ions; no molecules are left  Example: HCl + H 2 O  H 3 O + + Cl -  The Six Strong Acids: 1.HCl (hydrochloric acid) 2.HBr (hydrobromic acid) 3.HI (hydroiodic acid) 4.H 2 SO 4 (sulfuric acid) 5.HNO 3 (nitric acid) 6.HClO 4 (perchloric acid)

14. Strong acids:  HCl + H 2 O  H 3 O + + Cl- Remember, strong acids ionize completely in water. The reaction goes all the way to the right. A single arrow is used. There are virtually no HCl molecules left intact.

15. Strong acids:  HCl + H 2 O  H 3 O + + Cl- Looking at the equation above, with the single arrow, is Cl - a strong base or a weak base?

16. Weak Acids  weak acid – reacts only slightly with water to produce ions; mostly molecules left  HF + H 2 O  H 3 O + + F -  H 2 CO 3 + H 2 O  H 3 O + + HCO 3 -  Notice: Only one H comes off at a time

17. Strong Bases  Strong Base - reacts completely with water to produce ions; no molecules are left  Strong Bases:  Group 1 metals and Sr, Ba, and Ra with OH present  Examples NaOH, KOH, CsOH, Ba(OH) 2

18. Weak Bases  weak Bases – reacts only slightly with water to produce ions; mostly molecules left  Examples:  NH 3 + H 2 O   NH 4 + + OH -  Fe(OH) 3 + H 2 O   Fe 3+ + 3 OH -

19. General rule: The conjugate base of a strong acid is a weak base. Similarly, the conjugate acid of a strong base is a weak acid.

20. One more reminder: “Weak” does not mean the same thing as “diluted.” HCl, for example is always a strong acid. If you add 1000 liters of water to it, it will be diluted, but still strong because what little there is will be completely dissociated.

21. Concentrated Vs Dilute  Concentrated means that there is a lot of Acid/Base molecules for the amount of water or that the Molarity is High  Concentrated HCl is 16 M  Dilute means that there are not many Acid/Base molecules or that the Molarity is Low  Diluted HCl would be 0.1 M

22. Naming Acids Look at the names of these acids – can you come up with the rule? H 2 SO 4 : sulfuric acid HNO 3 : nitric acid H 3 PO 4 : phosphoric acid

23. Naming acids Rule #1: If the acid comes from a polyatomic ion that ends in “ate,” the acid is named ____-ic. H 2 SO 4 : sulfuric acid (from sulfate) HNO 3 : nitric acid (from nitrate) H 3 PO 4 : phosphoric acid (from phosphate)

24. Naming Acids Rule #2: If the acid does not have oxygen in it, then name it… hydro + second element + ic Example: HCl is hydrochloric acid. What would HBr be? H 2 S?

25. Water as an Acid and a Base  amphoteric – describes substance that can act as an acid or as a base  Example: H 2 O (see previous Bronsted Lowry examples)  Arrhenius:  H 2 O  H + + OH -  Bronsted Lowry:  H 2 O + H 2 O  H 3 O + + OH -

26. Strength of Acids/Bases  pH Scale relates strengths of acids and bases  pH 0 to 7 – Acid  pH = 7 – Neutral  pH 7 to 14 – Base  pH can only be from 0 to 14  pH= -log[H + ]

27. BasicAcidicNeutral 10 0 10 -1 10 -3 10 -5 10 -7 10 -9 10 -11 10 -13 10 -14 [H + ] 013579111314 pH Basic 10 0 10 -1 10 -3 10 -5 10 -7 10 -9 10 -11 10 -13 10 -14 [OH - ] 013579111314 pOH

28. pH  pH= -log[H + ]  Used because [H + ] is usually very small  As pH decreases, [H + ] increases exponentially  Sig figs only the digits after the decimal place of a pH are significant  [H + ] = 1.0 x 10 -8 pH= 8.00 3 sig figs

29. How to find pH  Punch in calculator: [H + ] or number, log, +/–  Example:  Find pH if [H + ] = 1.0 x 10 -5 M  pH = - log(1.0 x 10 -5 ) = 5.00  Punch in + / –, log, 1.0 EXP - 5, Enter

30. [H + ] Concentration  If you know the pH, you can determine the Molarity of H + Ions in the solution.  Since pH= -log [H + ]  Then [H + ] = 10 -pH  What is the Hydrogen Ion concentration of a solution that has a pH of 7.00?  [H + ] = 10 -7.00 = 1.0 x 10 -7 M

31. How to find [H + ]  Given pH, find [H]; Punch in pH or number, +/–, 10 x or 2nd log  Example: If pH is 9.0, find [H + ]  pH = -log [H + ]  Punch in 10 x or 2nd log, +/–, 9.0  [H] = 1.0 x 10 -9 M

32. What if it has OH - but not H + ?  14 = pH + pOH  1.0 x 10 -14 = [H + ] x [OH - ]  What is the pH of a solution that has a pOH of 5.00?  14 = pH + 5.00 or 14 – 5.00 = 9.00  What would the [H + ] be?  [H + ] = 10 -9.00 = 1.0 x 10 -9 M

33. pOH  Not all substances make [H + ], some make OH -  pOH= -log[OH - ]  14 = pH + pOH  Sig figs only the digits after the decimal place of a pH are significant  [OH - ] = 1.0 x 10 -8 pOH= 8.00 3 sig figs  pH = 14 – 8.00 = 6.00

34. Finding OH- if you have H+ 1.0 x 10-14 = [H+][OH-] 1.0 x 10-14 = [H+][OH-] If you hydrogen ion concentration is 2.50 x 10-5, what is the Hydroxide ion concentration? If you hydrogen ion concentration is 2.50 x 10-5, what is the Hydroxide ion concentration? 1.0 x 10-14 = [2.50 x 10-5][OH-] 1.0 x 10-14 = [2.50 x 10-5][OH-] 1.0 x 10-14 = [OH-] 1.0 x 10-14 = [OH-] 2.50 x 10-5= 4.0 x 10-10 2.50 x 10-5= 4.0 x 10-10

35. Indicators A solution or compound that changes color based on pH. Common Indicators: Universal Indicator (UI) Methyl Red Bromothymol Blue Methyl Orange