Acids & Bases

• Some are electrolytes Properties of Acids Sour taste Change color of acid-base indicators (red in pH paper) Some react with active metals to produce hydrogen gas Ba(s) + H2SO4(aq) BaSO4(s) + H2(g) Some react with bases to neutralize and form salt and water H2SO4 (aq) + 2NaOH(aq) Na2SO4 (aq) + 2H2O(l) • Some are electrolytes

Examples of Acids Lemons and oranges - citric acid Vinegar - 5% by mass acetic acid Pop and fertilizer - phosphoric acid

• Some are electrolytes Properties of Bases Bitter taste Change color of acid-base indicators (blue in pH paper) Dilute aqueous solutions feel slippery Ex. Soap • Some react with acids to neutralize and form salt and water • Some are electrolytes

Examples of Bases Soap - NaOH Household cleaners - NH3 Antacids - Ca(OH)2, Mg(OH)2

Arrhenius Acids Acids that increase the concentration of hydronium (H3O+) in aqueous solutions HNO3(aq) + H2O(l) H3O+(aq) + NO3-(aq) H+ + NO3- + H2O acid

Why do acids produce H3O+? H+ is extremely attracted to the unshared pair of electrons on the water molecule so it donates itself to this molecule where it becomes covalently bonded. The ion formed is known as the hydronium ion (H3O+) H+

NaOH(s) Na+(aq) + OH-(aq) Arrenius Bases Bases that increase the concentration of hydroxide ions (OH-) in aqueous solutions NaOH(s) Na+(aq) + OH-(aq) H2O

Strength of Acids & Bases Strong acids & bases completely ionize in aqueous solutions H2SO4 + H2O H3O+ + HSO4- NaOH Na+ + OH- Strong acids & bases are strong electrolytes A list of strong acids & bases can be found on pg. 460-461

Weak acids & bases only partially break down into ions when in aqueous solutions HCN + H2O H3O+ + CN- NH3 + H2O NH4+ + OH- • Weak acids & bases are weak electrolytes A list of weak acids & bases can be found on pg. 460-461

H2O(l) + H2O(l) H3O+(aq) + OH-(aq) Why can we drink H2O? Water self ionizes to form equal concentrations of H3O+ and OH- H2O(l) + H2O(l) H3O+(aq) + OH-(aq) A substance is considered “neutral” when [H3O+] = [OH-] [H3O+] concentration = 1.0 x 10-7M [OH-] concentration = 1.0 x 10-7 M

1.0 x 10-14 M2 = ionization constant for H2O (Kw) When [H3O+] = [OH-] If [H3O+] > 1.0 x 10-7 M, the solution is acidic If [OH-] > 1.0 x 10-7 M, the solution is basic To find the concentration of [H3O+] or [OH-] in acidic or basic solutions, the following equation can be used: 1.0 x 10-14 M2 = [H3O+] [OH-] 1.0 x 10-14 M2 = ionization constant for H2O (Kw)

Sample Problem A 1.0 x 10-4 M solution on HNO3 has been prepared for laboratory use. a. Calculate the [H3O+] of this solution b. Calculate the [OH-] of this solution c. Is this solution acidic or basic? Why? d. Substitute H2SO4 as the acid. How would the calculations change?

Sample Problem An aqueous 3.8 x 10-3 M NaOH solution has been prepared for laboratory use. a. Calculate the [H3O+] of this solution b. Calculate the [OH-] of this solution c. Is this solution acidic or basic? Why? d. Substitute Ca(OH)2 as the base. How would the calculations change?

Practice Problems Complete practice problems on pg. 484 #1-4

pH = -log [H3O+] [H3O+] = antilog (-pH) The pH scale The pH scale measures the power of the hydronium ion [H3O+] in a solution The scale typically goes from 1-14 (although it can extend below or above it under extreme conditions) The following equations can be used to determine the pH or [H3O+] of a solution: pH = -log [H3O+] [H3O+] = antilog (-pH) [H3O+] = 1 x 10-pH

pH > 7 basic pH = 7 neutral pH < 7 acidic

pOH = -log [OH-] [OH-] = antilog (-pOH) The pOH scale The pOH scale measures the power of the hydroxide ion [OH-] in a solution The scale typically goes from 1-14 (although it can extend below or above it under extreme conditions) The following equations can be used to determine the pOH or [OH-] of a solution: pOH = -log [OH-] [OH-] = antilog (-pOH) [OH-] = 1 x 10-pOH

pH + pOH = 14

Sample Problems Calculate the pH of each of the following. Classify as acidic or basic. 1.3 x 10-5 M NaOH 1.0 x 10-4 M HCl

Sample Problems What is the [H3O+] for each of the following? Classify as acidic or basic. pH = 5.8 b. pOH = 8.9

Sample Problems What is the [OH-] for each of the following? Classify as acidic or basic. [H3O+] = 9.5 x 10-10 M pOH = 1.3

Practice Problems Complete practice problems on pg. 487 #1

Strong Acid-Base Neutralization When equal parts of acid and base are present, neutralization occurs where a salt and water are formed HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

Sample Problems H2CO3 + Sr(OH)2 HClO4 + NaOH HBr + Ba(OH)2 NaHCO3 + H2SO4

Titrations When you have a solution with an unknown concentration, you can find it by reacting it completely with a solution of known concentration This process is known as titrating To perform a titration, an instrument called a buret can be used to precisely measure amounts of solution, drop by drop

Titration Termonology Equivalence point - the point at which the known and unknown concentration solutions are present in chemically equivalent amounts moles of acid = moles of base Indicator - a weak acid or base that is added to the solution with the unknown concentration before a titration so that it will change color or “indicate” when in a certain pH range (table 16-6 on pg. 495 in your text will show various indicators and their color ranges)

Phenolpthalein is clear at pH<8, End point - the point during a titration where an indicator changes color The 2 most common indicators we will use in our chemistry class will be: Phenolphthalein - turns very pale pink at a pH of 8-10 Bromothymol blue - turns pale green at a pH of 6.2-7.6 Phenolpthalein is clear at pH<8, pale pink at pH 8-10 and magenta at pH >10 Bromothymol blue

Practice Titration for an unknown acid 1. Titrate 5.0 of mL of unknown HCl into a 250 mL erlenmeyer flask - *remember to document the starting amount and ending amount of acid on the buret to prevent error 2. Add 2 drops of indicator (phenolphthalein) to the flask - the color of the solution should be clear 3. Titrate with .5M NaOH, continuously swirling the flask, until the solution turns very pale pink for 30 seconds - *remember to document the starting amount and ending amount of base on the buret 4. Mathematically determine the concentration of the unknown HCl solution by using the following equation:

Titration Equation MAVA = MBVB MA = molarity (mol/L) of acid VA = volume in L of acid MB = molarity (mol/L) of base VB = volume in L of base molesA = molesB 5. After calculating the molarity of the unknown acid experimentally, get the theoretical molarity and calculate % error

Practice titration for an unknown base 1. Titrate 5.0 of mL of unknown NaOH into a 250 mL erlenmeyer flask - *remember to document the starting amount and ending amount of base on the buret to prevent error 2. Add 2 drops of indicator (phenolphthalein) to the flask - the color of the solution should be magenta 3. Titrate with .5M HCl, continuously swirling the flask, until the solution turns very pale pink for 30 seconds - *remember to document the starting amount and ending amount of acid on the buret 4. Mathematically determine the concentration of the unknown NaOH solution by using MAVA = MBVB 5. After calculating the molarity of the unknown base experimentally, get the theoretical molarity and calculate % error