2. CHAPTER 16: (HOLT) ACID-BASE TITRATION AND pH

3. I. Concentration Units for Acids and Bases A. Chemical Equivalents 1. Definition: quantities of solutes that have equivalent combining capacity a. Acid: mass of one equivalent is numerically equal to the mass of one mole of the acid divided by the number of protons(H + or H 3 O + ) that one mole of the acid can provide Example: HCl 36 g/mol; 1 eq = 1H + ; 36 g/mol H + H 2 SO 4 98g/mol; 2 eq = 2H + ; 49 g/mol H +

4. B. Base: mass of one equivalent is numerically equal to the mass of one mole of the base divided by the number of protons(OH - ) that one mole of the base can provide Example: NaOH 40 g/mol; 1 eq = 1 OH - ; 40 g/mol OH - Ca(OH) 2 74 g/mol; 2 eq = 2 OH - ; 37 g/mol OH -

5. B. Normality Definition: number of equivalents of solute per liter of solution N = eq of solute L of solution

6. C. Relationship Between Normality and Molarity N = nM N: Normality n: number of equivalents (# of H + = or OH - ) M: Molarity Example: 1M HCl = 1N HCl 1M NaOH = 1N NaOH 1M H 2 SO 4 = 2N H 2 SO 4 1M Ca(OH) 2 = 2N Ca(OH) 2

7. II. Aqueous Solutions and the Concept of pH A. Self-Ionization of Water 1. Definition: Two water molecules interact to produce a hydronium ion and a hydroxide ion by proton transfer - forms a weak electrolyte 2. [ ] is symbol used to indicate concentration in moles per liter (Molarity) 3. H 2 O + H 2 O H 3 O + + OH - ; in pure water [H 3 O + ] = [OH - ] 4. [H 3 O + ][OH - ] = 10 -14 5. If the [H 3 O + ] increases then the [OH - ] decreases or If the [H 3 O + ] decreases then the [OH - ] increases

8. B. The pH scale 1. pH -- the negative of the common logarithm of the hydronium ion concentration pH = -log [H 3 O + ] 2. Acid: pH < 7 3. Base: pH > 7 4. Neutral: pH = 7

9. C. Calculations involving pH pH = -log [H 3 O + ] 0.001 M HCl = [H 3 O + ] =1 x 10 -3 pH = -log[1 x 10 -3 ] pH = 3 (acid) {Remember that [H 3 O + ][OH - ] = 1 x 10 -14 ; so if [H 3 O + ] = 1 x 10 -3 ; then [OH - ] = 1 x 10 -11 FYI: there is also pOH = - log[OH - ] and pH + pOH = 14

10. III. Acid-Base Titrations A. Indicators 1. Definitions: a. indicators - weak acid or base dyes whose colors are sensitive to pH, or hydronium, concentration b. transition interval - the pH range over which an indicator changes color

11. 2. Types of indicators a. Change color at about pH 7 b. Change color below pH 7 c. Change color above pH 7

12. B. The Principle of Titration Definitions: 1. Titration - the controlled addition and measurement of the amount of a solution of known concentration that is required to react completely with a measured amount of a solution of unknown concentration 2.Standard solution - a solution that contains a precisely known concentration of a solute

13. 3. Equivalence point - in a neutralization reaction, the point at which there are equivalent quantities of hydronium and hydroxide ions 4. End point - the point in a titration where an indicator changes color 5. Primary standard - a highly purified compound, when used in solution to check the concentration of the known solution in a titration

14. C. Molarity and Titration 1. Determine the moles of acid (or base) from the standard solution used during titration 2. From a balanced chemical equation, determine the ratio of moles of acid (base) to base (acid) 3. Determine the moles of solute of the unknown solution used during the titration 4. Determine the molarity of the unknown solution

15. D. Normality and Titrations V a x N a = V b x N b V a : volume of the acid N a : normality of the acid V b : volume of the base N b : normality of the base