Acids, Bases, and Salts

Objectives Know the fundamental properties of acids and bases. Know the ions associated with acids and bases. Know how acids and bases are produced.

Properties of Acids contain H (HCl, H2SO4, etc.) produce hydronium ions (H3O+) in water electrolytes pH < 7 sour taste react with metals to make H2 gas Zn + 2HCl → H2 + ZnCl2 made from non-metal oxide and water: SO3 + H2O →H2SO4

Properties of Bases often contain OH (NaOH, Ca(OH)2, etc.) bases produce hydroxide ions (OH–) in water electrolytes pH > 7 bitter taste feel slippery made from metal oxide and water: ZnO + H2O →Zn(OH)2 Acids and bases “neutralize” each other! HCl + NaOH → NaCl + H2O

Objectives Be able to name acids. Be able to identify Brønsted-Lowry acids, bases, conjugate acids, and conjugate bases. Understand and correctly apply the meaning of the term amphoteric.

Acid Nomenclature USE YOUR YELLOW SHEET! use the stem and ending of the anion name -ide hydro-stem-ic acid -ate stem-ic acid -ite stem-ous acid HCl = H+ + Cl– (chlor-ide) = hydrochloric acid HNO3 = H+ + NO3– (nitr-ate) = nitric acid HNO2 = H+ + NO2– (nitr-ite) = nitrous acid Common exceptions: sulfuric (H2SO4) and phosphoric (H3PO4)

Original Definitions of A/B Acids contain H, dissolve to form H+ Bases dissolve to form OH- HCl (g) → H+ (aq) + Cl− (aq) NaOH (s) → Na+ (aq) + OH− (aq) …but this is over-simplified.

Brønsted-Lowry Acids and Bases acid: proton (H+) donor base: proton (H+) acceptor HCl(g) + H2O(l) ↔ Cl−(aq) + H3O+(aq) hydronium ion acid base conjugate base conjugate acid NH3(aq) + H2O(l) ↔ NH4+(aq) + OH−(aq) conjugate acid conjugate base base acid HNO3(aq) + NH3(aq) ↔ NO3−(aq) + NH4+(aq) acid base conj base conj acid

Water: Acid and Base! amphoteric: a substance that can act as either an acid or a base (such as water) hydronium ion H3O+ = + ACID + BASE conj. acid BASE + hydroxide ion OH- = − ACID conj. base

Objectives Understand the process of self-ionization. Understand how the concentrations of hydronium and hydroxide ion can vary in water. Understand the concept of pH. Be able to make pH calculations using the log and 10x functions on a calculator.

Self-Ionization of Water H2O + H2O ↔ OH− + H3O+ (reactant strongly favored) [OH− ] = 10-7 M and [H3O+] = 10-7 M KW = [OH−]·[H3O+] water: Kw = [10-7]·[10-7] = 10-14 acidic: Kw = [10-9]·[10-5] = 10-14 basic: Kw = [10-3]·[10-11] = 10-14 [OH−] and [H3O+] are inversely proportional, KW = 10-14

[H3O+] ACIDS [H3O+] > 10-7 M memorize this! HNO3 (g) + H2O (l) →H3O+ (aq) + NO3− (aq) [H3O+] = 10-6 M, 10-5 M, 10-4 M, … 10-1 M or more BASES [H3O+] < 10-7 M memorize this! NaOH(s) → Na+(aq) + OH−(aq) OH− removes H3O+ [H3O+] = 10-8 M, 10-9 M, 10-10 M, … 10-14 M or less

pH Scale pH = −log[H3O+] [H3O+] = 10-3 M, pH = 3 (acidic) [H3O+] = 10-7 M, pH = 7 (neutral) [H3O+] = 10-11 M, pH = 11 (basic) Calculating pH? [H3O+] = 5.7 x 10-2 M pH = −log(5.7E-2) = 1.2 Calculating [H3O+]? Use [H3O+] = 10-pH If pH = 3.8 [H3O+] = 10-3.8 = 1.6 x 10-4 M

Objectives Understand the distinction between natural acid rain and human-caused (anthropogenic) acid rain. Know the types of pollution that result in acid rain formation. Understand the effects of acid rain and how they can be reduced.

Acid-Base Indicators ↔ compounds that change color at a specific pH phenolphthalein is a typical example… indicator anion indicator anion H+ ↔ H3O+ + + H2O ACID = clear CONJ BASE = pink add base (removes H3O+) = pink in high pH add acid (add H3O+) = clear in low pH transition occurs around pH = 9

Acid-Base Indicators many plants contain indicators: hydrangea, red cabbage universal indicator: mixture w/ wide pH range

Plant Dyes and pH serviceberry, willow bark, & Oregon grape root are local plants that contain indicators used as natural dyes for skins, feathers, etc. Coastal Salish

Objectives Understand the concept of KA and how it relates to strong and weak acids. Be able to calculate the KA of an acid solution if given the initial molarity and the pH of the solution.

Strengths of Acids strong acid: completely ionizes in water products favored (hydronium ions and acid anions) HNO3 (g) + H2O (l) → H3O+(aq) + NO3−(aq) weak acid: partially ionizes in water reactants favored (molecular form of acid) HC2H3O2(l) + H2O (l) ↔ H3O+(aq) + C2H3O2−(aq)

Acid Dissociation Constant (KA) HA + H2O ↔ H3O+ + A− (A is the acid anion) Strong acids—high KA ( > 1, products favored) Weak acids—low KA ( < 1, reactants favored) Note that [H3O+] = [A−] * use [H+] = 10-pH [HA] = C – [H+]

Calculating KA The initial concentration of an HNO2 solution is 0.315 M. What is the KA of HNO2 if the pH of the solution is 1.93? Determine [H3O+] (same value as [A-] ) [H3O+] = 10-pH = 10-1.93 = 0.012 M Determine [HA] [HA] = C – [H3O+] = 0.315 M – 0.012 M = 0.303 M Calculate KA KA < 1, weak acid

Objectives Be able to explain the distinction between strong and weak acids versus concentrated and dilute solutions. Understand the concept of acid neutralization and be able to determine the products of an acid-base neutralization reaction. Be able to calculate either acid or base concentration using data from an acid-base titration.

Strength vs. Concentration strength relates to degree of ionization (KA) concentration relates to amount of solute (molarity) strong = product favored (H+, A-) weak = reactant favored (HA) concentrated = lots of solute dilute = not much solute

Neutralization acid + base → “salt” + water H+ + OH− → H2O (neutral) salt: ionic compound consisting of a base cation and an acid anion HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (l) H2SO4 (aq) + 2KOH (aq) → K2SO4 (aq) + 2H2O (l) Try this one… HNO3 (aq) + Ca(OH)2 (aq) → ? + ? 2HNO3 (aq) + Ca(OH)2 (aq) → Ca(NO3)2 (aq) + 2H2O (l)

Acid-Base Titration standard solution (known concentration) is added to an unknown solution until pH = 7 the concentration of the unknown can be calculated

Titration Calculation What is the concentration of H2SO4 if 10.0 mL is completely neutralized by 14.2 mL of 1.0 M NaOH?

Buffers buffer: a solution in which the pH remains relatively constant when a small amount of acid or base is added consists of weak acid (or base) and one of its salts Example: Your blood pH (= 7.2) is maintained by H2CO3/HCO3− buffer Add acid: H+ + HCO3− → H2CO3 Add base: H2CO3 + OH− → HCO3− + H2O

Acid Rain normal rainfall is slightly acidic (pH = 5-6) CO2 + H2O → H2CO3 acid rain: precipitation with a pH < 5 burning “high-sulfur” coal produces SO2 & SO3 SO2 + H2O → H2SO3 cars make NOX that form HNO2 & HNO3

Impacts of Acid Rain destroys decomposes forests limestone kills aquatic life corrodes metal

Acid Rain in the USA

Neutralizing Acid Rain Limestone bedrock neutralizes acid, reducing environmental damage. Granite does not. Bases such as CaO or CaCO3 must be used to neutralize acids. H2SO4 + CaCO3 → CaSO4 + H2O + CO2