Acids and Bases

Acids and Bases Acid: when dissolved in water, increases the concentration of H+ (arrenhius acid) HCl H+ + Cl- HCl + H2O H3O+ + Cl- Strong acid: an acid that completely ionizes/dissociates in water HNO3 + H2O H3O+ + NO3- Strong electrolyte H2O

Bronsted-Lowry definition

Strong Acids: Hydrochloric Acid HCl Nitric Acid HNO3 Sulfuric Acid H2SO4 Perchloric Acid HClO4 Hydrobromic Acid HBr Hydroiodic Acid HI

Weak Acid: an acid that only partially ionizes/dissociates in water CH3COOH(aq) CH3COO-(aq) + H3O+(aq) Weak electrolytes Weak Acids: Phosphoric acid H3PO4 Acetic Acid CH3COOH Carbonic Acid H2CO3 Hydrocyanic Acid HCN Benzoic Acid C6H5COOH

Polyprotic Acids Polyprotic acids: acids that can release more than one H+ Sulfuric Acid H2SO4(aq) HSO4-(aq) + H3O+(aq) HSO4-(aq) SO42-(aq) + H3O+(aq)

Bases Base: a substance that, when put in water, increases the concentration of OH- ions or a substance that accepts H+ ions NaOH(aq)  Na+(aq) + OH-(aq) Strong Bases: bases that completely ionize in water

Weak Bases: bases that only partially ionize in water Ammonia: NH3 Pyridine: C5H5N

Each acid has a conjugate base and every base has a conjugate acid Congugate Acid-Base Pairs Each acid has a conjugate base and every base has a conjugate acid conjugate acid-base pair 1 HA + B A− + BH+ conjugate acid-base pair 2

Water is amphoteric – can act as an acid or a base

Neutralization Reactions When strong acids and bases in aqueous solution react with each other, they form water and a salt. HX(aq) + MOH(aq)  HOH(l) + MX(aq) HCl(aq) + NaOH(aq)  H2O(l) + NaCl(aq) Water Salt

Acid Ionization Constant Acid Ionization Constant (Ka): the equilibrium constant for the ionization reaction of an acid with water HA + H2O A- + H3O+ Large Ka = Strong acid Small Ka = Weak acid

Base Ionization Constant Base Ionization Constant (Kb): the equilibrium constant for the ionization reaction of a base with water B + H2O OH- + BH+ Large Kb = Strong base Small Kb = Weak base

Autoionization of Water

Autoionization of Water Also called “Self Ionization” About 1 out of every 10 million water molecules form ions through self ionization H2O Û H+ + OH– H2O + H2O Û H3O+ + OH– All aqueous solutions contain both H3O+ and OH–

Ion Product Constant for Water Ion Product Constant for Water (Kw): the numerical value obtained by multiplying the molar concentrations for hydronium and hydroxide ions present in pure water at 25°C Kw = [H3O+][OH-] = 1.00 x 10-14 at 25 oC the concentration of H3O+ and OH– are equal in pure water [H3O+] = [OH–] = 10-7M @ 25°C

Ion Product of Water the product of the H3O+ and OH– concentrations is always the same number Kw =[H3O+][OH–] = 1.00 x 10-14 @ 25°C if you measure one of the concentrations, you can calculate the other as [H3O+] increases the [OH–] must decrease so the product stays constant inversely proportional

H+ OH- [H+] vs. [OH-] Acid Base Even though it may look like it, neither H+ nor OH- will ever be 0

Acidic and Basic Solutions Neutral solutions have equal [H3O+] and [OH–] [H3O+] = [OH–] = 1 x 10-7 acidic solutions have a larger [H3O+] than [OH–] [H3O+] > [OH–] [H3O+] > 1 x 10-7; [OH–] < 1 x 10-7 basic solutions have a larger [OH–] than [H3O+] [H3O+] < [OH–] [H3O+] < 1 x 10-7; [OH–] > 1 x 10-7

pH and pOH Acidic solutions Basic solutions Neutral solutions

pH + pOH = 14 pH = -log [H+] OR pH = -log [H3O+] [H3O+] = 10-pH pH is a measure of the concentration of H3O+ in solution pH = -log [H+] OR pH = -log [H3O+] [H3O+] = 10-pH pOH is a measure of the concentration of OH- in solution pOH = -log [OH-] [OH-] = 10-pOH pH + pOH = 14

Classification of Water Soluble Substances Electrolytes: solutes that separate into ions when dissolved in water (they’re soluble) Have the ability to conduct electricity 2 types Strong electrolytes Weak electrolytes

Strong electrolytes: solutes that completely dissociates into ions when dissolved in water Examples: NaCl, MgBr2, HCl Strong electrical conductors Strong electrolyte(aq or s) → Cation+(aq) + Anion-(aq) Example: NaCl(s) → Na+(aq) + Cl-(aq)

Examples: HF, NH3, acetic acid Weak electrical conductors Weak Electrolytes: solutes that, when dissolved in water, only partially dissociates into ions Examples: HF, NH3, acetic acid Weak electrical conductors Weak electrolyte(aq) ↔ Cation+(aq) + Anion-(aq) Example: HF(aq) ↔ H+(aq) + F-(aq)

Nonelectrolyte (s or l) → Nonelectrolyte(aq) Nonelectrolytes: solutes that dissolve in water without separating into ions Examples: sucrose, ethanol Do not conduct electricity Nonelectrolyte (s or l) → Nonelectrolyte(aq) Example: C12H22O11(s) → C12H22O11(aq)

Titration Titration: a procedure for the quantitative analysis of a substance of unknown concentration whereby a measured quantity of another substance, of know concentration, is completely reacted with the with the original substance. Often used to determine the concentration of acids and bases

Equivalence point: the point in a titration at which one reactant has been exactly consumed by the addition of another reactant Midpoint of vertical rise Occurs at pH = 7 in a strong acid-strong base titration [H3O+] = [OH-]

Indicators Acid-Base Indicator: a chemical that changes color with a change in pH Added to solutions in small amounts in order to determine to solution’s pH visually Usually organic compounds Weak acid or base establishes an equilibrium with the H2O and H3O+ in the solution 42

HInd(aq) + H2O(l)  Ind(aq) + H3O+(aq)

Phenolphthalein 44 44

Bromocresol Green Yellow Green Blue

Methyl Red 46