Acids and Bases Chapter 19

Review Electrolyte A substance that conducts an electrical current when melted or in solution Ionic compounds Acids and Bases

Acid-Base Theories Different definitions of acids and bases Arrhenius Bronsted-Lowry

Arrhenius Acid Compounds that ionize to produce hydrogen ions (H+) in aqueous solutions Examples: HCl, HBr, H2SO4, CH3COOH *Note: CH3COOH is an organic acid Acidic H

Arrhenius Base Compounds that ionize to produce hydroxide ions (OH-) in aqueous solutions Examples: KOH, NaOH, LiOH *Note: CH3OH is not a base, it’s an organic alcohol

Bronsted-Lowry (alternate) Acid Hydrogen ion donor Examples: HCl, HBr, H3O+ Base Hydrogen ion acceptor Examples: H2O, NH3

Acid-Base Theories Bronsted- Lowry Arrhenius

Properties of Acids & Bases Acids - Taste Sour Bases - Taste Bitter, Feel Slippery Will change color of acid – base indicator Can be strong or weak electrolytes in an aqueous solution

Ionization Electrolytes will dissociate into ions when dissolved in water Strong Electrolytes will completely dissociate Weak Electrolytes will only partially dissociate

Ionization HCl(s)  H+(aq) + Cl-(aq) HNO3(s)  H+(aq) + NO3-(aq) NaOH(s)  Na+(aq) + OH-(aq) KOH(s)  K+(aq) + OH-(aq)

Polyprotic Acids Acids that have more than one H Can release more than one H+ into solution Examples:H2SO4, H3PO4 H2SO4(s)  2H+(aq) + SO42-(aq) Bases can also release more than one OH- into solution Mg(OH)2(s)  Mg+2(aq) + 2OH-(aq)

Ionization of Water Water can be split into 2 ions Ionization of Water H+ and OH- Ionization of Water H2O  H+ + OH- H2O + H2O  H3O+ + OH-

Ionization of Water H+(aq) = H3O+

Neutral Solutions For neutral solutions For all aqueous solutions [H+] = [OH-] For all aqueous solutions [H+] * [OH-] = 1.0 x 10-14 [ ] means concentration

Measuring Acidity (Alkalinity) Traditionally we measure [H+] pH = -log [H+] Neutral solution [H+] = 1.0 x 10-7 pH = 7 pOH = -log [OH-] pH + pOH = 14

Acidity Acidic Solutions pH < 7.0 Basic Solutions pH > 7.0 [H+] > 1.0 x 10-7 Basic Solutions pH > 7.0 [H+] < 1.0 x 10-7

Measuring pH Litmus paper pH paper pH Meter Red in acid Blue in base pH paper pH Meter Acid – Base Indicators (Table M)

Table M

Changes in pH pH increases by 1 for every decrease in [H+] by a magnitude of 10 [H+] pH 1.0*10-7 7 1.0*10-8 8 1.0*10-9 9 1.0*10-10 10

pH Changes

pH Changes

Neutralization Acid + Base  Water + Salt HA + BOH  HOH + BA Double Replacement Reaction HA + BOH  HOH + BA

Neutralization Examples: HCl + NaOH  H2O + NaCl HNO3 + LiOH  H2O + LiNO3 H2SO4 + 2KOH  2 H2O + K2SO4

Titration Process in which a volume of solution known concentration is used to determine the concentration of another solution Usually shown by a color change of an indicator (end point)

Titration Example (6M)(1L) = (X)(2L) X =3M NaOH 2 Liters of an unknown conc. of NaOH is titrated with 1 Liter of 6M HCl, what is the concentration of the base? MAVA = MBVB (6M)(1L) = (X)(2L) X =3M NaOH

Polyprotic Acids Acids that have more than one H [H2SO4] = 2M Examples:H2SO4, H3PO4 [H2SO4] = 2M [H+] = 4M [H3PO4] = 2M [H+] = 6M

Titration (Reality) MAVA(#of H’s) = MBVB(#of OH’s)

Another Titration Example 500 milliliters of an unknown conc. of NaOH is titrated with 1 Liter of 1M H2SO4, what is the concentration of the base? MAVA = MBVB (1M)(1L)(2) = (X)(0.5L) X =4M NaOH