1. Unit 14 Acids, Bases and Salts

2. Operational Definitions: those that are observable in the lab Acids: Aqueous solutions of acids conduct electricity (because ions form H+, H 3 O+) Acids react with certain metals to form hydrogen gas ex: 2HCl(aq) + Zn(s) = ZnCl 2 (aq) + H 2 (g) *Table J* Any metal higher on the list will replace H

3. Acids cause acid-base indicators to change color –Litmus paper turns red –Phenolphthalein remains colorless Acids react with bases to form salt and water ex: HCl + NaOH = NaCl + H 2 O Dilute aqueous solutions of acids taste sour. Responsible for the sour taste of fruit and vinegar.

4. Bases As with acids, aqueous solutions of bases conduct electricity because ions form OH-_ Bases cause acid-base indicators to change color –Litmus paper turns blue –Phenolphthalein hot pink

5. Bases are caustic/corrosive, they can cause chemical burns to skin. Bases emulsify oils. This means they can dissolve oils and can dissolve in water and as a result they make good soaps. Bases taste bitter Bases feel slippery

6. Conceptual Definitions of Acids and Bases Those that cannot be directly observed. These are concepts that are based on other observations

7. Arrhenius theory of acids and bases Acids produce H+ as the only positive ion in solution. In water H+ forms a coordinate covalent bond with the water forming H 3 O+ (hydronium ion) ex: HCl → H+ + Cl- H 2 SO 4 → 2H + + SO 4 -2

8. The Nature of the Hydrogen Ion Hydrogen consists of one electron and one proton. When hydrogen becomes a positive ion it has lost its electron and is just a proton. *A positive hydrogen ion is a proton.* This proton cannot exist by itself, so it bonds with water to form a hydronium ion.

9. Acids that contain one H are called monoprotic; acids that contain two H are called diprotic; acids that contain three H are called triprotic. Polyprotic is another term used to describe acids that contain more than one H. Example: HCl – monoprotic H 2 SO 4 - diprotic H 3 PO 4 - triprotic

10. Bases produce OH- as the only negative ion in solution

11. The Nature of the Hydroxide Ion The presence of the hydroxide ion makes that base an electrolyte. It also gives the base its properties of being slippery and bitter to the taste. Ammonia is a base because when dissolved in water produces an OH ion.

12. Bases that contain one OH are called monohydroxy; acids that contain two OH are called dihydroxy. Example: NaOH – monohydroxy Ca(OH) 2 - dihydroxy

13. An alternative to Arrhenius Theory Hydrogen ion donors and acceptors (Bronsted-Lowry Theory) Acids donate hydrogen ion (donate a proton) Bases accept hydrogen ion (accept a proton)

14. *More general than Arrhenius’ definition and they do not require aqueous solutions, so water is not necessarily a product. Example: NH 3 + H 2 O → NH 4 + OH- HCO 3 - + H 2 O → CO 3 -2 + H 3 O +

15. *Water is one of several “special” substances. In certain reactions it reacts like an acid (H+ donor) while in other reactions it reacts like a base (H+ acceptor). These substances are known as amphiproticor amphoteric substances.

16. Reference Tables *Common Acids can be found on Table K (Strongest being the top 3) *Common Bases can be found on Table L (Strongest being the top 2) *Common Acid-Base indicators are found on Table M

17. Ionization the reaction between solute and solvent resulting in the formation of ions. All acids and bases will ionize in solution. The degree to which they ionize will determine the strength of the acid or base. This will also determine how good or poor an electrolyte the substance is. Salts will also ionize when placed in water because they are ionic substances in which H+ is not the only + ion in solution and OH- is not the only – ion in solution. Example: KCl → K + + Cl-

18. Dissociation Results in almost complete separation of ions in which very little molecule is left as a result. Very strong acids and bases dissociate in water. Example: HCl → H + + Cl-

19. Ionization Constants (k) are used to determine the strength of an acid or base based on its ability to produce H+ or OH-. The larger the k, the more ions produced, the stronger the acid (base) and better the electrolytic abilities of the solution.

20. Neutralization Reactions The reaction between an acid and a base to produce a salt and water Example: nitric acid + potassium hydroxide → salt + water Now, put it into a formula: HNO 3 + KOH → H 2 O + KNO 3

21. Titration a process in which the strength of an unknown acid (base) is determined by dropping in a known base (acid) until an endpoint (near neutral) is reached.

22. Acid + Base → Salt + Water For this reaction to become completely neutral, the number of moles of acid ion (H+) must be equal to the number of moles of base ion (OH-). *Formula (found on Table T) M A V A = M B V B *M A = molarity of H+ in the acid M B = molarity of OH- in the base V A = volume of the acid V B = volume of the base

23. Examples –What is the molarity of NaOH if 5mL are needed to exactly neutralize 10mL of 0.1M HCl? M A V A = M B V B (.1)(10) = (x)(5) x =.2M NaOH

24. More than 1 H or OH –What is the molarity of 25mL of H 2 SO 4 if 25mL of 2M KOH completely neutralize the solution? *(2)M A V A = M B V B *2(H) vs 1(OH) (2)(x)(25) = (2)(25) 50x = 50 x = 1M H 2 SO 4

25. Salts Salts are ionic substances that result from the reaction between an acid and a base. Some common examples of salts are NaCl, KNO 3, Ca(ClO 3 ) 2 …

26. Salts can be slightly acidic, slightly basic, or neutral in nature depending upon the acids and bases that are used in producing them.

27. pH and the pH scaleThis scale involves negative logarithms and powers of 10 pH is the percentage of hydrogen ion concentration in an aqueous solution

28. pH can be considered percent H+ or H 3 O+ Excess H+ in an aqueous solution will produce H 3 O+ which is a result of a coordinate covalent bond between the H+ and H 2 O pOH determines the percentage of OH-

29. pH scale the pH scale is usually shown as numbers between 1 and 14 1714 Acid Base Acids will yield H+ or H 3 O+ in solution. Bases will yield OH- in solution.

30. determine the pH (or pOH) of a solution [H+] = 1 x 10 -pH [H+][OH-] = 1 x 10 –14 Questions: What is the hydroxide ion concentration of a solution with a pH of 3? pH = 3 [H+] = 1x10 -3 [OH-] = 1x10 -11

31. What is the pH of a solution with a hydroxide ion concentration of 1 x 10 -13 ? [OH-] = 1x10 -13 [H+] = 1x10 -1 pH=1

32. Given a.001M solution of HCl, what is the pH? 1x10 -3 acid [H+] = 1x10 -3 pH=3

33. Given a.01M NaOH solution, what is the pH? 1x10 -2 base [OH-] = 1x10 -2 so [H+} = 1x10 -12 pH=12