Chemistry Chapter 20-21 Acids and Bases

Electrolytes: Solutions that conduct electrical current easily. Aqueous solutions of acids and bases are electrolytes. Ionic solids that dissolve in water are electrolytes.

Electrolytes: The ability of a solution to conduct an electric current. The ammeter measures the flow of electrons (current) through the circuit. If the ammeter measures a current, and the bulb glows, then the solution conducts. (Electrolyte) If the ammeter fails to measure a current, and the bulb does not glow, the solution is non-conducting. (Non-electrolyte)

NaCl(s)  Na+(aq) + Cl-(aq) Dissociation: Water surrounds ions and dissolves them. Dissociation equation: NaCl(s)  Na+(aq) + Cl-(aq)

Dissociation of HCl HCl  H+ + Cl- HA  H+ + A- equation. HCl  H+ + Cl- General Dissociation equation. HA  H+ + A- A- represents the anion

Basic definition: Acids: Produce hydrogen ions, H+ , when dissolved in water. Bases: Produce hydroxide ions, OH-, when dissolved in water.

Acid-Base Neutralization + Base  Water + Salt HCl + NaOH  HOH H2O + NaCl

Some Properties of Acids Acids taste sour Acids react with certain metals to produce hydrogen gas. 2 HCl(aq) + Mg(s)  H2(g) + MgCl2(aq) Acids effect indicators Blue litmus turns red Phenolphthalein turns colorless Methyl orange turns red

Acids Effect on Indicators Blue litmus paper turns red in contact with an acid.

Bases Effect on Indicators Red litmus paper turns blue in contact with a base. Phenolphthalein turns purple in a base.

pH Indicators and their ranges

Hydro_______ic acid Naming Acids Acids without oxygen: HF = Hydrofluoric acid HCl = Hydrochloric acid HBr = Hydrobromic acid HI = Hydroiodic acid H2S = Hydrosulfuric acid HCN = Hydrocyanic acid CN- = cyanide anion

Naming Acids H+ + (anions with oxygen) Acids with oxygen: “ate-ic ite-ous” Anion name Acid name ____ate ____ic acid ___ite ___ous acid HNO3 = Nitr__?__ acid Anion NO3-1 = nitrate Nitric acid HNO2 = Nitr__?__ acid Anion NO2-1 = nitrite Nitrous acid

Naming Acids H+ + (anions with oxygen) Formula H+ with Anion Acid name H2SO4 2H+ + SO42- = sulfate Sulfuric acid H2SO3 2H+ + SO32- = sulfite Sulfurous acid HC2H3O2 H+ + C2H3O2-1 = Acetate Acetic acid H2CO3 2H+ + CO32- = Carbonate Carbonic acid H3PO4 3H+ + PO43- = Phosphate Phosphoric acid

Naming Acids H+ + (series of Cl and O) Formula H+ with Anion Acid name HCl H+ + Cl- = Chloride Hydrochloric acid HClO H+ + ClO-1 = hypochlorite Hypochlorous acid HClO2 H+ + ClO2-1 = Chlorite Chlorous acid HClO3 H+ + ClO3-1 = Chlorate Chloric acid HClO4 H+ + ClO4-1 = perchlorate Perchloric acid Hypo = less oxygen Per = more oxygen

Naming Bases Same method as naming ionic compounds Formula Cation with OH- Base name NaOH Na+ + OH- Sodium hydroxide (Lye) KOH K+ + OH- Potassium hydroxide Ca(OH)2 Ca2+ + 2 OH- Calcium hydroxide (Lime) Mg(OH)2 Mg2+ + 2 OH- Magnesium hydroxide (Milk of magnesia) Al(OH)3 Al3+ + 3 OH- Aluminum hydroxide NH3 (No ions) Ammonia (common name) Molecules that contain nitrogen are weak bases.

(Self-Ionization of Water) Two water molecules collide to form Hydronium and Hydroxide ions. H2O + H2O  H3O+ + OH- 1 in 10,000,000 water molecules self-ionize. This is 10-7 water molecules.

Kw – Ionization Constant for Water In pure water at 25 C: [H+] = [H3O+] = 1 x 10-7 mol/L 1 10,000,000 (Dissociate) [OH-] = 1 x 10-7 mol/L At 25 C, 1 in 10,000,000 water molecules will dissociate and form H+ and OH-. Kw = Ion product constant (for water at 25 C) Kw = [H+] x [OH-] Kw = (1 x 10-7)(1 x 10-7) = 1 x 10-14

Kw = [H+] x [OH-] Ion-product constant = Hydrogen-ion concentration x hydroxide-ion concentration

Ion Concentration in Water

Calculating pH “per Hydronium” pH = -log10[H3O+] In other words: pH = -log[H+] [H+] = H+ concentration in (mol/L) or (M) Example: [H+] = 1 x 10-7 M pH = -log(1 x 10-7) log(1) = 0 log (10-7) = -7 pH = -(0-7) = 7.0

Calculating pH Examples: [H+] = 1 x 10-1 pH = -log(1 x 10-1)

Acids Have a pH less than 7

Bases have a pH greater than 7

Kw = [H+] x [OH-] Ion-product constant = Hydrogen-ion concentration x hydroxide-ion concentration Basic, neutral and acidic are all ways of describing pH concentrations in a solution.

Calculating pH Examples: [H+] = 1 x 10-1 pH = 1.0 pH = 4.0 Acid pH = 4.0 Acid [H+] = 1 x 10-4 pH = 7.0 Neutral [H+] = 1 x 10-7 pH = 9.0 Base [H+] = 1 x 10-9 pH = 13.0 Base [H+] = 1 x 10-13 pH = 0 Acid [H+] = 1 x 10 0

Calculating pH Examples: pH + pOH = 14 [OH-] = 1 x 10-2 pOH = 2.0 Base 14.0 [OH-] = 1 x 10-5 pOH = 5.0 pH = 9.0 Base 14.0 [OH-] = 1 x 10-7 pOH = 7.0 pH = 7.0 Neutral 14.0

Calculating pH pH = -log [H+] pOH = -log [OH-] pH + pOH = 14 [OH-] = 1 x 10-8 pOH = 8.0 pH = 6.0 Acid 14.0 [OH-] = 1 x 10-10 pOH = 10.0 pH = 4.0 Acid 14.0 [OH-] = 1 x 10-14 pOH = 14.0 pH = 0 Acid 14.0

Calculating pH pH = -log [H+] pOH = -log [OH-] pH + pOH = 14 [OH-] = 1 x 10 0 pOH = 0 pH = 14.0 Base 14.0

Examples (Using Calculator): pH = -log [H+] pOH = -log [OH-] pH + pOH = 14 Examples (Using Calculator): [H+] = 2.5 x 10-5 pH = -log(2.5 x 10-5) pH = 4.60 Acid [H+] = 3.0 x 10-9 pH = -log(3.0 x 10-9) pH = 8.52 Base

pH = -log [H+] pOH = -log [OH-] pH + pOH = 14 Examples: [OH-] = 6.0 x 10-3 pOH = -log(6.0 x 10-3) pOH = 2.22 pH = 11.78 Base 14.0 [OH-] = 9.0 x 10-12 pOH = -log(9.0 x 10-12) pOH = 11.05 pH = 2.95 Acid 14.0

Equation Summary pH = -log [H+] pOH = -log [OH-] pH + pOH = 14 Kw = [H+] x [OH-] [OH-] = 10-pOH 1 x 10-14 = [H+] x [OH-] [H+] = 10-pH pH + pOH = 14 [OH-] = 10-pOH pOH = 14 - pH [OH-] = 10-9 [H+] = 10-5 Acid pOH = 14 - 5 9 1 x 10-9 M 5 1 x 10-5 M pOH = ___ pH = ___ [H+] = __________ [OH-] = __________ pH + pOH = 14 [H+] = 10-pH [OH-] = 10-pOH pH = 14 - pOH [OH-] = 10-6 [H+] = 10-8 pH = 14 - 6 6 8 pOH = ___ 1 x 10-6 M pH = ___ [H+] = __________ 1 x 10-8 M [OH-] = __________ Base

Equation Summary pH = -log [H+] pOH = -log [OH-] pH + pOH = 14 Kw = [H+] x [OH-] [OH-] = 10-pOH 1 x 10-14 = [H+] x [OH-] [H+] = 10-pH pH + pOH = 14 [OH-] = 10-pOH pOH = 14 - pH [OH-] = 10-11.2 [H+] = 10-2.8 Acid pOH = 14 – 2.8 [H+] = 0.0016 11.2 6.3 x 10-12 M 2.8 1.6 x 10-3 M pOH = ___ pH = ___ [H+] = __________ [OH-] = __________

Acid-Base Theories Theory Acid Base Arrhenius Produces H+ ions in solution Produces OH- ions in solution HCl HC2H3O2 NaOH Ca(OH)2

Acid-Base Theories Base Acid NH3 + HOH  HOH NH4+ + OH- Acid Base Theory Acid Base Bronsted-Lowry Proton (H+ ion) donor. Proton (H+ ion) acceptor. Includes more molecules that do not have OH- NH3 (ammonia) Conjugate Acid-Base Pairs Base (proton acceptor) Acid (Conjugate acid) proton donor NH3 + HOH  HOH NH4+ + OH- Acid (proton donor) Base (conjugate base) proton acceptor

Ionization of HCl and formation of hydronium ion, H3O+ H2O + HCl  H3O+ + Cl- Proton acceptor Proton donor Hydronium ion Chloride ion Acid Conjugate Base Bronsted/Lowry (Acid-Base Pair)

Ionization of HCl and formation of hydronium ion, H3O+ H2O + HCl  H3O+ + Cl- Proton acceptor Proton donor Hydronium ion Chloride ion Conjugate Acid Base Another Acid-Base Pair

Acids are Proton Donors Bronsted/Lowry Theory = proton donor idea Monoprotic acids Diprotic acids Triprotic acids H3PO4 HCl H2SO4 HNO3 H2CO3 HC2H3O2 Note: a proton is a hydrogen ion. “protic” = proton (All theories agree these are acids)

Acid-Base Theories Theory Acid Base Lewis (e- pair) acceptor (e- pair) donor BCl3 NH3 B N H Cl Cl H H Cl

Lewis Acid-Base Theory Lewis Base e- pair acceptor e- pair donor Cl Al Cl N H Cl H H AlCl3 (aluminum chloride) NH3 (ammonia)

Lewis Acid-Base Theory Lewis Base e- pair acceptor e- pair donor Al Cl N H AlCl3 (aluminum chloride) NH3 (ammonia)

Acid-Base Theories Summary Theory Acid Base Arrhenius Produces H+ ions in solution Produces OH- ions in solution HCl HC2H3O2 NaOH Ca(OH)2 HOH HOH Bronsted-Lowry Amphoteric: A substance that can act as an acid or a base. Water is an example: H2O = HOH H+ ion (proton) donor. H+ ion (proton) acceptor. NH3 (ammonia) Lewis (e- pair) acceptor (e- pair) donor NH3 BCl3

Strong Acid Dissociation H+ + H2O  H3O+ Hydronium is a result of proton transfer (H+ donated to H2O).

Strong Acids Strong acids are assumed to be 100% ionized in solution (good proton donors). HCl H+ + Cl- H2SO4 H+ + HSO4- HNO3 H+ + NO3- Arrow in one direction indicates that molecule dissociates completely.

Weak Acids Weak acids are usually less than 5% ionized in solution (poor proton donors). HF H+ + F- H3PO4 H+ + H2PO4- H2PO4- H+ + HPO42- Double arrow = (Reaction goes both ways) The larger arrow = (Favored side)

Weak Acid Dissociation Weak acids are mostly molecules. (Only 5% or less dissociate)

Acid Dissociation Constant (Weak Acids) Ka = product concentrations vs. reactant concentrations for weak acids. [HF] HF(aq) + H2O(l) [H3O+] H3O+(aq) + F-(aq) [F-] Ka = H2O not included because it is in a different phase.

Base Dissociation Constant (Weak Bases) Kb = product concentrations vs. reactant concentrations for weak acids. [NH3] NH3(aq) + HOH(l) [NH4+] NH4+(aq) + OH-(aq) [OH-] Kb = H2O not included because it is in a different phase.

HX(aq) H+(aq) + X-(aq) [H+] [X-] Ka = = = [HX] Initial Change 0.35 M Change 0.041 M 0.041 M - 0.041 M Equilibrium 0.35 – 0.041 = 0.309 M 0.041 M 0.041 M [H+] [X-] (0.041) (0.041) Ka = = = 0.0054 M [HX] 0.309

Methyl Red changes from red to orange in this range. Phenolphthalein changes to pink in this range.

HIn(aq) H+(aq) + In-(aq)

Phenolphthalein  Bromthymol blue  Methyl Red (Weakest acid) (Strongest Acid)

H2In(aq)  H+(aq) + HIn-(aq) Red to yellow. The indicator must be a diprotic acid for 2 color changes. H2In(aq)  H+(aq) + HIn-(aq) Red to yellow. HIn-(aq)  H+(aq) + In2-(aq) Yellow to blue.

The other terms are all associated with bases. Kw = [H+] x [OH-] Ion-product constant = Hydrogen-ion concentration x hydroxide-ion concentration Basic, neutral and acidic are all ways of describing pH concentrations in a solution. The other terms are all associated with bases. The other terms are related to pH or can be calculated from pH.

The other terms are all acid-base theories. Conjugate acids and bases are part of Bronsted-Lowry theory. The other terms describe ionizable hydrogen ions.. Dissociation constants are used with weak acids and bases.

Sulfuric Acid Used in the production of paper Highest volume production of any chemical in the U.S. Used in the production of paper Used in production of fertilizers Used in petroleum refining

Nitric Acid Used in the production of fertilizers Used in the production of explosives Nitric acid is a volatile acid – its reactive components evaporate easily Stains proteins (including skin!)

Hydrochloric Acid Used in the pickling of steel Used to purify magnesium from sea water Part of gastric juice, it aids in the digestion of protein Sold commercially as “Muriatic acid”

Phosphoric Acid A flavoring agent in sodas Used in the manufacture of detergents Used in the manufacture of fertilizers Not a common laboratory reagent

Acetic Acid Used in the manufacture of plastics Used in making pharmaceuticals Acetic acid is the acid present in vinegar

Examples of Organic Acids Citric acid in citrus fruit Malic acid in sour apples Deoxyribonucleic acid, DNA Amino acids, the building blocks of protein Lactic acid in sour milk and sore muscles Butyric acid in rancid butter

Weak Acids: (Organic) Organic acids (Acids that contain carbon) are always weak acids. HC2H3O2 H+ + C2H3O2- CH3COOH H+ + CH3COO- Same acid written another way Carboxylic Acid group: R-COOH

Organic Acids Organic acids all contain the “carboxyl” group called “carboxylic acid”, sometimes several of them. The carboxyl group is a poor proton donor, so ALL organic acids are weak acids.

Properties of Bases Bases effect indicators Red litmus turns blue Bases taste bitter Bases effect indicators Red litmus turns blue Phenolphthalein turns purple Bases have a pH greater than 7 Bases are proton (hydrogen ion, H+) acceptors Solutions of bases feel slippery Bases neutralize acids

Products of Neutralization HCl + NaOH  NaCl + H2O H2SO4 + Ca(OH)2  CaSO4 + 2 H2O HNO3 + KOH  KNO3 + H2O The products of neutralization are always a ______ and _______. salt water

Measuring pH with wide-range paper

Narrow-Range pH Paper

pH Indicators and their ranges

Calculating [H+] and [OH-] [H+] = 10-pH and [OH-] = 10-pOH Examples: pH = 2.0 [H+] = 1 x 10-2 Acidic pH = 12.0 Basic [H+] = 10-12 pH = 4.3 Acidic [H+] = 10-4.3 = 5 x 10-5 pOH = 4.0 [OH-] = 1 x 10-4 pH=10 Basic pOH = 8.0 [OH-] = 1 x 10-8 pH=6 Acidic pOH = 10.0 [OH-] = 1 x 10-10 pH=4 Acidic

SUMMARY Calculating pH, pOH pH = -log[H3O+] or pH = -log[H+] pOH = -log[OH-] Relationship between pH and pOH pH + pOH = 14 Finding [H3O+], [OH-] from pH, pOH [H3O+] = 10-pH or [H+] = 10-pH [OH-] = 10-pOH

Acids React with Active Metals Acids react with active metals to form salts and hydrogen gas. Mg + 2HCl  MgCl2 + H2(g)

Acids React with Carbonates 2HC2H3O2 + Na2CO3 2 NaC2H3O2 + H2O + CO2

Effects of Acid Rain on Marble (calcium carbonate) George Washington: BEFORE George Washington: AFTER

Acids Neutralize Bases HCl + NaOH  NaCl + H2O Neutralization reactions ALWAYS produce a salt and water.

Products of Neutralization (Lab 23) Trial 1 (Strong acid and Strong Base) HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l) Trial 2 (Weak acid and Strong Base) HC2H3O2(aq) + NaOH(aq)  NaC2H3O2 (aq) + H2O(l) Trial 3 (Strong acid and Weak Base) HCl(aq) + NH3 (aq)  NH4+1(aq) + Cl-1(aq) Trial 4 (Weak acid and Weak Base) HC2H3O2(aq) + NH3(aq)  NH4+1(aq) + C2H3O2-1(aq)

Titration Curves

Bases Neutralize Acids Milk of Magnesia contains magnesium hydroxide, Mg(OH)2, which neutralizes stomach acid, HCl. 2 HCl + Mg(OH)2 MgCl2 + 2 H2O

Titration Calculations