1. ACIDS & BASES

2. ACID/BASE THEORY Acids and bases are solutions which can be described differently by multiple theories. So far, we have treated everything that utilizes “H” as a cation = acid H 3 PO 4 “OH” as an anion = base NaOH

3. ARRHENIUS THEORY The theory that anything capable of producing H + ions are acids and those producing OH - anions are bases is Arrhenius’ Theory of acids/bases. Based on the notion that these substances dissociate in solution into their ions. H 2 SO 4 2H + + SO 4 -2 H20H20

4. ACID OR BASE? HNO 3 KOH AcidBase Ca(OH) 2 HBr BaseAcid HC 2 H 3 O 2 Al(OH) 3 AcidBase

5. BRONSTAD LOWRY ACIDS/BASES More general description of acids and bases, incorporates all Arrhenius acid/bases Acids = H + donors Bases = H + acceptors Example: NH 3 (aq) + H 2 O (l) NH 4 + (aq) + OH - (aq) Would solutions of NH 3 be acidic or basic?

6. CONJUGATE ACIDS/BASES NH 3 (aq) + H 2 O (l) NH 4 + (aq) + OH - (aq) Conjugate Acid: The compound remaining after H + acceptance by a base. NH 3 NH 4 + Base Conjugate Acid Conjugate Base: The compound remaining after H + donation by an acid. H 2 O OH - Acid Conjugate Base

7. LEWIS ACID/BASES Describes Acid/Bases based upon electron donation. Lewis Acid: A substance that can accept a pair of electrons to form a covalent bond. Lewis Base: A substance that can donate a pair of electrons to form a covalent bond. H + + OH - H 2 O

8. STRENGTH OF ACIDS/BASES For this class, we will mostly discuss Arrhenius acid/bases (H + cation/OH - anion) The strength of these solutions is determined by the concentration (Molarity) of H + or OH - ions in the solution. Dependent upon… 1) Amount of dissolved solute donating the ion 2) Likelihood of ion dissociation

9. CONCENTRATION OF SOLUTE Just like calculating the molarity of non acid/base solutions. M = mol solute / liter solution Example: What is the concentration of a H 2 SO 4 solution if 13.2 g of solute are dissolved in 100mL of di-water?

10. CALCULATING [ H + ] & [OH - ] However, we cannot assume the molarity (concentration) of H + /OH - ions is the same as the molarity of the solution for 2 reasons: 1.Not all acids/bases completely dissociate into their ions. 2.Not all acidic/basic compounds donate 1 ion per molecule

11. Complete dissociation: all solute dissociates into ions, meaning the concentration of ions can be calculated with the molar ratio. HClH + + Cl - 1 mol HCl produces 1 mol H + So a 0.2M HCl solution would have a H + concentration of 0.2M as well DISSOCIATION

12. COMPLETE DISSOCIATION Few acids/bases completely dissociate If they do, they are termed “Strong Acids/Bases” because they produce high concentrations of their ions upon dissociation. Strong Acids (7): HCl, HBr, HI, H 2 SO 4, HNO 3, HClO 3, HClO 4 Strong Bases (8): LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH) 2, Sr(OH) 2, Ba(OH) 2

13. CALCULATING ION DONATION Not all acids/bases donate 1 H + or OH - into solution when they dissociate! How many ions would the following acids/bases donate if they completely dissociate? HBr Ca(OH) 2 H 2 SO 4

14. [H+] EXAMPLE Using the previous H 2 SO 4 example (13.2 g of solute are dissolved in 100mL of di- water), calculate the [H 3 O + ] of the solution.

15. PARTIAL DISSOCIATION Weak acids/bases do not completely dissociate in water. They are called weak, because a lower concentration of H + or OH - ions results. Example: Only about 1% of Phosphoric Acid (H 3 PO 4 ) dissociates into H + + H 2 PO 4.

16. DISSOCIATION OF STRONG & WEAK ACIDS/BASES In other words, you would need to know what percentage of an acidic/basic compound is likely to donate its ions into solution in order to calculate the [H + ] & [OH - ]. For this class, we will be working with strong acids/bases. So complete dissociation (ion donation) can be assumed.

17. POTENTIAL OF HYDROGEN pH = potential of hydrogen pH is an expression of the concentration of H + ions in solution. Since H + readily reacts with water to form H 3 O +, we calculate… pH = −log [H 3 O + ] [H 3 O + ] = 10 -pH

18. PH & POH Just as pH is an expression of the concentration of H + ions in solution, pOH expresses the concentration of OH- ions in solution… pOH = −log [OH - ] [OH - ] = 10 -pOH pH & pOH always add up to 14 and thus are related through the equation: pH + pOH = 14

19. PH SCALE pH < 7 acidic pH = 7 neutral pH > 7 basic High [H 3 O + ] acid Low [H 3 O + ] base [H 3 O + ] = ~1x10 -7 neutral

20. ACID/BASE REACTIONS Also known as neutralization reactions, because the pH of the final solution becomes neutralized towards pH=7 Either [H 3 O + ] or [OH - ] is decreased by forming more neutral products. Acid + Base Salt + Water

21. ACID/BASE REACTIONS 2 Types you will need to know… 1) Acid + Base Salt + H 2 O 2) Acid + Carbonate Salt + H 2 O + CO 2

22. ACID/BASE REACTION #1 Example: Hydrochloric acid reacts with sodium hydroxide. HCl(aq) + NaOH(aq) NaCl(aq) + H 2 O(l)

23. ACID/BASE REACTION #2 Example: Potassium carbonate reacts with sulfuric acid. H 2 SO 4 (aq)+ K 2 CO 3 (aq) K 2 SO 4 (aq) + H 2 O(l) + CO 2 (g) Acid Carbonate Salt Water Carbon Dioxide