Acids and Bases ACIDS and BASES “In nature, acids can be found in fruits: citric acid is responsible for the sharp taste of lemons. Vinegar contains acetic acid, and tannic acid from tree bark is used to tan leather. The stronger mineral acids have been prepared since the Middle Ages. One of these, aqua fortis (nitric acid), was used by assayers to separate gold from silver. Car batteries contain sulfuric acid, also strong and corrosive. A base is the opposite of an acid. Bases often feel slippery; bicarbonate of soda and soap are bases, and so is lye, a substance that can burn skin. Bases that dissolve in water are called alkalis. In water, acids produce hydroxide ions. When an acid and a base react together, the hydrogen and hydroxide ions combine and neutralize each other, forming water together and a salt. The strength of acids and bases can be measured on a pH scale.” Eyewitness Science “Chemistry” , Dr. Ann Newmark, DK Publishing, Inc., 1993, pg 42

Terms Acids Examples: Bases Compounds that form H+ in water. Formulas usually begin with ‘H’. Examples: HCl – hydrochloric acid HNO3 – nitric acid H2SO4 – sulfuric acid Bases Compounds that form OH- in water Formulas usually end in ‘hydroxide’

Properties ACIDS BASES electrolytes electrolytes sour taste bitter taste turn litmus red turn litmus blue react with metals to form H2 gas slippery feel vinegar, milk, soda, apples, citrus fruits ammonia, lye, antacid, baking soda ChemASAP

Acid + Base  Salt + Water Orange juice + milk  bad taste Evergreen shrub + concrete  dead bush HCl + NaOH  NaCl + HOH salt water

Common Acids and Bases Strong Acids (strong electrolytes) Strong Bases (strong electrolytes) NaOH sodium hydroxide KOH potassium hydroxide Ca(OH)2 calcium hydroxide HCl hydrochloric acid HNO3 nitric acid HClO4 perchloric acid H2SO4 sulfuric acid Phosphoric acid, H3PO4, is another acid commonly found in the laboratory. It is on the borderline between a strong and weak acid. Weak Base (weak electrolyte) NH4OH ammonia Weak Acids (weak electrolytes) CH3COOH acetic acid H2CO3 carbonic Weak Base (weak electrolyte) NH3 ammonia NH3 + H2O  NH4OH Kotz, Purcell, Chemistry & Chemical Reactivity 1991, page 145

Acid Nomenclature A binary acid – is an acid that contains only two different elements: hydrogen and one of the more electronegative elements. Rules of naming a binary acid: The name of the binary acid begins with the prefix hydro- The root of the name of the second element follows this prefix. The name then ends with the suffix –ic A oxyacid – is an acid that is a compound of hydrogen, oxygen, and a third element, usually a nonmetal.

Acid Nomenclature

Acid Nomenclature HBr H2CO3  hydrobromic acid H2SO3  carbonic acid 2 elements, -ide  hydrobromic acid 3 elements, -ate  carbonic acid 3 elements, -ite  sulfurous acid

Acid Nomenclature hydrofluoric acid sulfuric acid nitrous acid  H+ F- 2 elements  H+ F-  HF 3 elements, -ic  H+ SO42-  H2SO4 3 elements, -ous  H+ NO2-  HNO2

Common Industrial Acids Sulfuric Acid Petroleum refining Metallurgy Manufacture of manure Essential to the processes of: production of metals, paper, paint, dyes, detergents and in automobile batteries. Effective dehydration agent. Nitric Acid Is a volatile, unstable liquid rarely used in industry. Dissolving in water provides stability. Used in the making of explosives, rubber, plastics, dyes and pharmaceuticals. Stains proteins yellow.

Common Industrial Acids Phosphoric Acid Used directly for manufacturing of fertilizers and animal feed, detergents and ceramics. Flavoring of beverages. Cleaning agent for dairy equipment. Hydrochloric Acid Aids in digestion. Used for “pickling” iron and steel. General cleaning agent. Used for food processing, activation of oil wells, recovery of Mg from sea water, and the production of other chemicals.

Common Industrial Acids Acetic Acid Used for synthesizing chemicals used in the manufacture of plastics. Used as a fungicide. Food supplements

pH of Common Products

: measures acidity/basicity pH scale : measures acidity/basicity ACID BASE 10x 100x 10x 10x 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 NEUTRAL Each step on pH scale represents a factor of 10. pH 5 vs. pH 6 (10X more acidic) pH 3 vs. pH 5 (100X different) pH 8 vs. pH 13 (100,000X different)

: measures acidity/basicity pH scale : measures acidity/basicity Soren Sorensen (1868 - 1939) ACID BASE 10x 100x 10x 10x 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 NEUTRAL Each step on pH scale represents a factor of 10. pH 5 vs. pH 6 (10X more acidic) pH 3 vs. pH 5 (100X different) pH 8 vs. pH 13 (100,000X different)

Kw = [H3O+][OH-] = 1.0  10-14 H2O + H2O H3O+ + OH- Ionization of Water H2O + H2O H3O+ + OH- Kw = [H3O+][OH-] = 1.0  10-14

Ionization of Water Find the hydroxide ion concentration of 3.0  10-2 M HCl. [H3O+][OH-] = 1.0  10-14 [3.0  10-2][OH-] = 1.0  10-14 [OH-] = 3.3  10-13 M Acidic or basic? Acidic

The pH Scale pH – of a sol’n is the negative of the common logarithm of the hydronium ion concentration [H3O+]. pOH - of a sol’n is the negative of the common logarithm of the hydroxide ion concentration [OH-].

pouvoir hydrogène (Fr.) pH Scale 14 7 INCREASING ACIDITY INCREASING BASICITY NEUTRAL pH = -log[H3O+] pouvoir hydrogène (Fr.) “hydrogen power”

pH = -log[H3O+] pOH = -log[OH-] pH + pOH = 14 pH Scale pH = -log[H3O+] pOH = -log[OH-] pH + pOH = 14

pH Scale What is the pH of 0.050 M HNO3? pH = -log[H3O+] Acidic or basic? Acidic

pH Scale What is the molarity of HBr in a solution that has a pOH of 9.6? pH + pOH = 14 pH + 9.6 = 14 pH = 4.4 pH = -log[H3O+] [H3O+] = 10-pH [H3O+] = 10-4.4 [H3O+] = 4.0  10-5 M HBr Acidic

The pH Scale Basic pH and pOH calculations: pH = -log[H3O+] pOH = -log[OH-] pH = 14 – pOH pOH = 14 – pH Calculating pH from [H3O+] Determine concentration of [H3O+] using Kw formula. Determine pH using logarithm formula. Calculating pOH from [OH-] Determine concentration of [OH-] using Kw formula. Determine pOH using logarithm formula.

Example pH (pOH) calculations If you had a sol’n with the concentration of 0.025 M [H3O+], what is the sol’ns pH? What is the sol’ns pOH? 1. Given: [H3O+] = 0.025 M Unknown: pH = ?? pH = -log [H3O+] = -log[0.025M] = 1.6 pOH = 14 – pH = 14 - 1.6 = 12.4

Example pH (pOH) calculations B. What is the pH of 0.050 M HNO3? Is it acidic or basic? pH = -log[H3O+] pH = -log[0.050] = 1.3 It is acidic because 1.3  7.

Example pH (pOH) calculations C. What is the pH of a 0.100 M HCl sol’n? Given: Identity Concentration: [0.100 M] Unknown: pH = ???? 2. [H3O+] = 0.100 M pH = -log[0.100 M] = 1.00

Example pH (pOH) calculations D. What is the pH of a 1.5 M sol’n of KOH? 1. Given: Identity Concentration: [1.5 M] Unknown: pH = ???? 2. [OH-] = 1.5 M [H3O+] = 1.0 x 10-14 M2/[OH-] = 1.0 x 10-14 M2/[1.5 M] = 6.7 x 10-15M pH = -log[H3O+] = -log[6.7 x 10-15M] = 14.17 Check : Answer does indicate that KOH forms a sol’n pH  7, which is basic.

Example pH (pOH) calculations E. Determine the hydronium concentration and the hydroxide concentration of an aqueous sol’n that has a pH of 5.6. 1. Given: pH = 5.6 Unknown: [H3O+] = ???? [OH-] = ???? 2. [H3O+] = 10-pH = 10-5.6 = 2.5 x 10-6 M [OH-] = 1.0 x 10-14 M2/[H3O+] = 1.0 x 10-14 M2/[2.6 x 10-6] = 3.8 x 10-9 M Check: The pH is acidic, therefore it should be found that the [H3O+]  [OH-]. Yes it is.

Example pH (pOH) calculations F. What is the molarity of HBr in a solution that has a pOH of 9.6? 1. Given: pOH = 9.6 Unknown: pH = ???? [H3O+] = ???? 2. pH + pOH = 14 pH = 14 – 9.6 = 4.4 [H3O+] = 10-pH = 10-4.4 = 4.0 x 10-5 M

The Three Acid/Base Theories Arrhenius Bronsted-Lowry Lewis

A Comparison b/w acid-base theories Type Acid Base Arrhenius H+ or H3O+ producer OH- producer Bronsted-Lowry Proton donor (H+) Proton acceptor (H+) Lewis Electron-pair acceptor Electron-pair donor

HCl + H2O  H3O+ + Cl– – + B. Definitions Arrhenius - In aqueous solution… Acids form hydronium ions (H3O+) HCl + H2O  H3O+ + Cl– H Cl O – + acid

NH3 + H2O  NH4+ + OH- – + B. Definitions Arrhenius - In aqueous solution… Bases form hydroxide ions (OH-) NH3 + H2O  NH4+ + OH- H N O – + base

HCl + H2O  Cl– + H3O+ B. Definitions acid conjugate base base Brønsted-Lowry Acids are proton (H+) donors. Bases are proton (H+) acceptors. HCl + H2O  Cl– + H3O+ acid conjugate base base conjugate acid

B. Definitions H2O + HNO3  H3O+ + NO3– B A CA CB

Amphoteric - can be an acid or a base. B. Definitions NH3 + H2O  NH4+ + OH- B A CA CB Amphoteric - can be an acid or a base.

Polyprotic - an acid with more than one H+ B. Definitions Give the conjugate base for each of the following: HF H3PO4 H3O+ F - H2PO4- H2O Polyprotic - an acid with more than one H+

Br - HSO4- CO32- HBr H2SO4 HCO3- B. Definitions Give the conjugate acid for each of the following: Br - HSO4- CO32- HBr H2SO4 HCO3-

Range and Color Changes of Some Common Acid-Base Indicators pH Scale 1 2 3 4 5 6 7 8 9 10 11 12 13 14 Indicators Methyl orange red 3.1 – 4.4 yellow Methyl red red 4.4 6.2 yellow Bromothymol blue yellow 6.2 7.6 blue Neutral red red 6.8 8.0 yellow From F. Brescia et al., Chemistry: A Modern Introduction, W. B. Saunders Co., 1978. Adapted from R. Bates, Determination of pH, Theory and Practice, John Wiley & Sons, Inc., New York, 1964. Phenolphthalein colorless 8.0 10.0 red colorless beyond 13.0

Common pH Indicators Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 520

Edible Acid-Base Indicators COLOR CHANGES AS A FUNCTION OF pH INDICATOR pH 2 3 4 5 6 7 8 9 10 11 12 RED APPLE SKIN BEETS BLUEBERRIES RED CABBAGE CHERRIES GRAPE JUICE RED ONION YELLOW ONION PEACH SKIN PEAR SKIN PLUM SKIN RADISH SKIN RHUBARB SKIN TOMATO TURNIP SKIN * Source: Volume 62, Number 4, April 1985 pg 285 (Not sure of magazine title) *YELLOW at pH 12 and above

Phenolphthalein Indicator Colorless = Acidic pH Pink = Basic pH

Acid-Base Neutralization 1+ 1- + + Hydronium ion Hydroxide ion H3O+ OH- Water H2O Water H2O Dorin, Demmin, Gabel, Chemistry The Study of Matter 3rd Edition, page 584

Acid-Base Neutralization 1+ 1- + + H3O+ OH- H2O H2O Hydronium ion Hydroxide ion Water Water Dorin, Demmin, Gabel, Chemistry The Study of Matter 3rd Edition, page 584

Titration Acids and bases sometimes react in a 1:1 mole ratio. HCL(aq) + NaOH(aq)→ NaCl(aq) + H2O(l) When sulfuric acid and sodium hydroxide react…the ratio is 1:2. H2SO4(aq) + 2NaOH(aq) → Na2SO4(aq) + 2H2O(l) This acid - base mole ratio from the balanced equation is a conversion factor used when neutralizing an acid or a base

Titration – using mole ratios… How many moles of sulfuric acid would you require to neutralize 0.50 mol of sodium hydroxide? STEP 1: Write the balanced equation for the rxn to find acid:base mol ratio H2SO4(aq) + 2NaOH(aq) → Na2SO4(aq) + 2H2O(l) STEP 2: Use the mole ratio to calculate the necessary number of moles of H2SO4 0.50 mol NaOH x 1 mol H2SO4 = 0.25 mol H2SO4 2 mol NaOH

Titration – Your Turn  How many moles of sodium hydroxide are required to neutralize 0.20 mol of nitric acid? STEP 1: NaOH + HNO3 → NaNO3 + H2O STEP 2: 0.20 mol HNO3 x 1 mol NaOH 1 mol HNO3 ANSWER: 0.20 mol NaOH needed How many moles of potassium hydroxide are needed to neutralize 1.56 mol of phosphoric acid?

Acid-Base Titration

1: 30 mL of 0. 10M NaOH neutralized 25. 0mL of hydrochloric acid 1: 30 mL of 0.10M NaOH neutralized 25.0mL of hydrochloric acid. Determine the concentration of the acid Write the balanced chemical equation for the reaction Extract the relevant information from the question: NaOH      V = 30mL , M = 0.10M       HCl    V = 25.0mL, M = ? Calculate moles NaOH       n(NaOH) = M x V = From the balanced chemical equation find the mole ratio  NaOH:HCl       Find moles HCl Calculate concentration of HCl using the formula: M=n÷V

2. What volume of 0. 75M H2SO4 is required to neutralize 25. 0 mL of 0 2. What volume of 0.75M H2SO4 is required to neutralize 25.0 mL of 0.427M KOH? Write the balanced chemical equation for the reaction Extract the relevant information from the question: KOH      V = 25.0mL , M = 0.427M       HCl    V = ?mL, M= 0.75M Calculate moles KOH       n(KOH) = M x V = From the balanced chemical equation find the mole ratio  KOH: H2SO4       Find moles H2SO4 Calculate volume of H2SO4 using the formula: V=n÷M

3. 50. 0 mL of an unknown solution of Ca(OH)2 are titrated with 0 3. 50.0 mL of an unknown solution of Ca(OH)2 are titrated with 0.15M HCl. Find the molarity of the Ca(OH)2 solution if 83 mL of acid are required to reach the equivalence point. Write the balanced chemical equation for the reaction Extract the relevant information from the question: Ca(OH)2   V = 50.0mL , M = ?M       HCl    V = 83 mL, M= 0.15M Calculate moles HCl       n(HCl) = M x V = From the balanced chemical equation find the mole ratio  Ca(OH)2 : HCl       Find moles Ca(OH)2 Calculate molarity of Ca(OH)2 using the formula: M=n÷V

Titration Curve

Titration – Virtual Lab http://lrs.ed.uiuc.edu/students/mihyewon/chemlab_experiment.html http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/stoichiometry/acid_base.html