Unit 12: Acids, Bases, and Salts

Properties of Acids Taste sour Are electrolytes: substances that produce ions in solution and as a result conduct electricity Strong acid  strong electrolyte Weak acid  weak electrolyte Will react with active metals (above H2) to produce H2(g)

Review of Naming Rules: from Unit 5! Binary Acids: hydro_______ic acid Ternary Acids (use Table E) ate ic ite ous

Properties of Bases Strong bases consist of a Group 1 or Group 2 metal w/ hydroxide ion Taste bitter, feel slippery Also electrolytes Strong bases  strong electrolytes Weak bases  weak electrolytes Will NOT REACT with active metals

Look on RT for common acids and bases

Electrolytes Form ions in water Conduct electricity Salts Ionic Compounds M to NM Acids pH < 7 H+ > OH- Bases pH > 7 OH- > H+ Molecules (NM to NM) are NOT electrolytes…do NOT produce ions in solution!!

Indicators- Table M Below the range  the color is given on the left Above the range  the color is given on the right.

Arrhenius Theory of Acids and Bases Arrhenius Acid Arrhenius Base A substance, that when dissolved in water, produces H+ ions as the only positive ions in solution A substance, that when dissolved in water, produces OH- ions as the only negative ions in solution Monoprotic, diprotic, triprotic Hydronium ion HCl (aq)  H+ (aq)+ Cl- (aq) NaOH (aq)  Na+ (aq) + OH- (aq) known as a hydrogen ion known as a hydroxide ion

Bronsted-Lowry Theory of Acids and Bases ***based on the proton (H+) Bronsted Acid Bronsted Base A proton donor A proton acceptor Must release an H+ (aq) ion Does not need to have the OH-(aq) ion, but MUST have a lone pair of electrons. Ex) HCl (aq) Ex) NH3 (aq)

Reversible Acid-Base Rxns Acid-Base reactions can be reversible! H2SO4(aq) + NH3(aq) HSO4-(aq) + NH4+(aq) B.A. B.B. B.B. B.A. A conjugate acid-base pair consists of two substances related by the loss or gain of a single hydrogen ion (H+) Strong acids have weak conjugate bases Strong bases have weak conjugate acids Pg.9

Amphoteric Substances A substance that can act as both an acid and a base is said to be amphoteric Ex) HCO3- (aq) + H2O (l)  CO32- (aq) + H3O+(aq) **Here, HCO3- acts as an acid Ex) HCO3- (aq) + H2O (l)  H2CO3 (aq) + OH- (aq) **Here, HCO3- acts as a base Water is a common amphoteric substance- can act as both an acid and base

Lewis Definition of Acids and Bases Lewis Acid Lewis Base an electron pair ACCEPTOR An electron pair DONOR Ex) NH3 + HCl  NH4+ + Cl-

Strengths of Acids and Bases STRONG acids/bases dissociate completely in solution HCl(aq) H+(aq) + Cl-(aq) Ka = [H+] [Cl-] = very large [HCl] STRONG acids have large Ka’s

Strength of acids and bases CH3COOH(aq) CH3COO-(aq) + H+(aq) Ka = [CH3COO-] [H+] = very small (less than 1) [CH3COOH] Not a lot of H+, therefore, weak acid weak acids have small Ka’s

MEMORIZE STRONG ACIDS Weak acids HI HBr HCl HNO3 H2SO4 HClO4 CH3COOH (vinegar) H3PO4 HF Strong acids have weak conjugate bases

MEMORIZE STRONG BASES Weak bases Any group 1 or group 2 metal with OH- NH3 Strong bases have weak conjugate acids

For diprotic and triprotic acids, the H+’s come off in steps For diprotic and triprotic acids, the H+’s come off in steps. It is more difficult to remove subsequent H+’s. Ex. Ka H3PO4  H+ + H2PO4- 10-3 H2PO4-  H+ + HPO4-2 10-8 HPO42-  H+ + PO4-3 10-13

Water as an Acid/Base Water molecules undergo self-ionization H2O(l) + H2O (l) H3O+ (aq) + OH- (aq) In pure water, the equilibrium concentration of [H3O+] = [OH-] = 1 x 10-7 M In other words, a neutral solution has the same number of moles of H+ ions as OH- ions.

“See-Saw Diagram” Ion-Product Constant for Water (for aqueous solutions) Kw= [H+] x [OH-] = 1.0 x 10-14 ** exponents always add to 14

The pH scale Scale is based on the concentration of H+ ions in aqueous solution pH = - log [H+] pH < 7, acidic pH = 7, neutral pH > 7, basic

Determining pH given concentrations [H+] = 1.0 x 10-6 M pH = 6 pH + pOH ALWAYS equals 14 [OH-] = 1.0 x 10-8 M pOH = 8

The pH scale is logarithmic An increase in one unit of the pH scale represents a ten-fold decrease (x 10) in the H+ ion concentration

Practice pH 6 to pH 4 pH 5 to pH 6 [H+] increases by a factor of 100 [H+] decreases by a factor of 10 pH 2 to pH 4 [H+] decreases by factor of 100

pH of weak acids Q. What is the pH of 0.10 M CH3COOH? CH3COOH(aq)  H+(aq) + CH3COO-(aq) 0.10 M – x x x Ka= [H+] [CH3COO-] = X2 = 1.8 x 10-5 [CH3COOH] 0.10-x x2 = 1.8 x 10-5 0.10-x x2= 1.8 x 10-5(0.10) x2= 1.8 x 10-6 x= 1.3 x 10-3 pH= -log(1.3 x 10-3)= 2.9 Very small. Ignore if 5% or less dissociation 0.10 M HCl pH=1 0.10 M CH3COOH pH=2.9

Neutralization Reactions Neutralization occurs when an acid and a base react to form a salt and water Acid + Base  salt + water HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l) Acid Base Salt Water Net Rxn: H+ (aq) + OH- (aq)  H2O (l)

Titrations Used to determine the concentration of an acid or base The equivalence point is reached when the moles H+ = moles OH- Titration Equation: MaVa = MbVb ** Can use any unit of volume as long as you are consistent

unknown volume, unknown concentration Known volume, known concentration

Indicators for Titrations Phenolphthalein is used to approximate the endpoint of a titration. Colorless indicates an acidic solution and pink indicates a basic solution When a pale pink color persists, the endpoint has been reached The endpoint, or equivalence point, is when moles H+ = moles OH- (pH =7)

Warm-Up What is the pH of a 0.025 M solution of Ca(OH)2? What is the hydroxide ion concentration of a 0.0001 M solution of HNO3 (aq)? What is the pH of a 0.220 M solution of hydrofluoric acid? (Assume 3% dissociation)

Hydrolysis (reverse of neutralization) The process in which ions of a salt react with water to produce an acidic or basic solution is called hydrolysis

Examples Example Parent Base Parent Acid NaCl NaOH HCl Neutral  strong strong NaCH3COO NaOH CH3COOH Basic  strong weak NH4NO3 NH4OH HNO3 Acidic weak strong NH4NO2 NH4OH HNO2 pH? Depends on Ka/Kb values weak weak

Salt formed from pH SB SA neutral SB WA More than 7 WB SA Less than 7 WB WA ? Need Ka/Kb

Basic Anhydrides Group I and II metal oxides that can react with water to form basic solutions. Ex. Na2O(s) + H2O(l)  2NaOH(aq) Basic anhydride

Acid Anhydrides Non-metal oxides that can react with water to form acidic solutions Produce acid rain Ex. SO3(g) + H2O(l)  H2SO4(aq) Ex. N2O5(g) + H2O(l)  2HNO3(aq) Ex. CO2(g) + H2O(l)  H2CO3(aq)

Chemistry of Dyeing Eggs https://www.youtube.com/watch?v=lMQ_qhEe yl8