1. Acids and Bases Chp 16

2. Old Definitions  Classic –Acids taste sour –Bases taste bitter  Arrhenius model –Acids produce hydronium ions (H 3 O + ) in solution –Bases produce hydroxide ions (OH - ) in solution

3. Current Definition  Bronsted – Lowry Model –Acids donate protons –Bases accept protons

4. Conjugate Acid/Base Pairs  When an acid donates a proton, it forms a conjugate base –HCl donates and becomes Cl -  When a base accepts a proton, it forms a conjugate acid –OH - accepts and becomes H(OH)  These differ by only ONE hydrogen atom  If the acid is strong, it’s conjugate base is weak (and vice versa)

5. Acid-Base Reactions  Can occur in either direction, but both sides compete for the free H ions so 1 direction often dominates HCl + H 2 O H 3 O + + Cl - Acid base conj. acid conj. base

6. Some Terms  Amphoteric –can be acids or bases, depending on the circumstance –Would be weak in both cases –Water is the most common amphoteric substance  Buffer –A weak acid or base that can be added to a solution to help it resist changes in pH

7. Indicators  Substances that change color when exposed to an acid or a base  Occurs because of reactions with the hydrogen ions  Many substances are natural pH indicators (ex. Blueberries, red cabbage, ammonia)  Some are commonly used in labs –Litmus paper (bases turn red paper blue, acids turn blue paper red) –Phenalthalien turns red when basic

8. Concentrations  Can be measured with molarity, but can be misleading because some acids have multiple H ions to donate  Use Normality instead: N = mols (equivalents) Liters Liters  Equivalents = how many H ions (or OH ions) that acid or base has  Ex. 0.5 mol H 3 PO 4 dissolved in 1.5 L has what normality? N = 0.5(3)=1 N 1.5 1.5

9. Neutralization  Adding just enough acid to completely react with a base (or vice versa)  Need to have every proton used  Creates water and a salt (ionic compound)  Can determine the amount of acid or base in a reaction using this in a process known as titration –Adding a little of a known acid or base until the equivalence point is reached as shown by a pH indicator  N acid V acid = N base V base

10. An Example  What volume of 0.10 M NaOH is required to completely neutralize 50.0 mL of 0.20 M H 2 SO 4 ? 0.10 M NaOH = 0.10 (1) = 0.10 N NaOH 0.20 M H 2 SO 4 = 0.20 (2) = 0.40 N H 2 SO 4 SO 0.40 (50) = 0.10 (V) V = 200 ml

11. [H+] and [OH-] Ions  Acids create H + ions when mixed in water because they donate H + to it  Bases create OH - ions when mixed in water because they take a H + away from it  If [H + ] > [OH - ], the solution is acidic  If [H + ] < [OH - ], the solution is basic  If [H + ] = [OH - ], the solution is neutral  However, for any solution [H + ] [OH - ] = 1 x 10 -14

12. An example  If a solution has a [OH - ] = 3.4 x 10 -9 M, what is the concentration of [H + ]? [H + ] (3.4 x 10 -9 ) = 1 x 10 -14 [H + ] (3.4 x 10 -9 ) = 1 x 10 -14 [H + ] = 2.9 x 10 -6 M [H + ] = 2.9 x 10 -6 M  Is the solution acid or basic? since [H + ] > [OH - ], it is acidic (remember that negative exponents mean more zeros after the decimal, so a smaller #)

13. pH Scale  An easier way to measure acids and bases because the concentrations we work with are so small  0-7 acidic  7 neutral  7-14 basic  The closer to 7 the weaker the substance is, the further away the stronger it is  Each step is actually a factor of 10 –So pH 2 is 100 X stronger than pH 4

14. pH continued  pH measures the concentration of H ions (or hydronium ions) in solution  pH = -log [H + ]  pOH = -log [OH - ]  pH + pOH = 14  To get [H + ] from pH: [H + ] = inverse log (-pH) [H + ] = inverse log (-pH)

15. An example  What is the pH of a solution that has a [H + ] of 1.3 x 10 -11 M? pH = -log (1.3 x 10 -11 ) = 10.88  Is it an acid or a base? basic since the pH is greater than 7

16. Another Example  What is the [H + ] if the pH of a solution is 2.3? [H + ] = inverse log (-2.3) = 5.01 x 10 -3 M [H + ] = inverse log (-2.3) = 5.01 x 10 -3 M  What would the pOH of this solution be? 2.3 + pOH =14 pOH = 11.7