Acids, Bases, and Salts

Objectives Know the fundamental properties of acids and bases. Be able to identify an Arrhenius acid. Be able to write a dissociation equation for an Arrhenius acid. Be able to identify monoprotic, diprotic, and triprotic acids.

Properties of Acids ACIDS pH < 7 sour taste electrolytes react with metals to make hydrogen gas Zn + 2HCl → H2 + ZnCl2 often formed from non-metal oxide and water SO3 + H2O →H2SO4

Properties of Bases BASES pH > 7 bitter taste electrolytes feel slippery often formed from metal oxide and water ZnO + H2O →Zn(OH)2 Acids and bases “neutralize” each other! HCl + NaOH → NaCl + H2O

Arrhenius Acids What explains acidic properties? Arrhenius acid: contains H and produces H+ ions in water Acids are molecules that ionize in water. HCl(g) →H+(aq) + Cl–(aq) H2SO4(l) →H+(aq) + HSO4–(aq) monoprotic: HC2H3O2 diprotic: H2CO3 triprotic: H3PO4 Svante Arrhenius 1859-1927

Objectives Be able to name acids. Be able to identify an Arrhenius base and write a dissociation equation for an Arrhenius base. 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)

Arrhenius Bases Bases dissociate to form OH- (hydroxide) ions when aqueous. NaOH(s) → Na+(aq) + OH-(aq) Mg(OH)2(s) → Mg2+(aq) + 2OH-(aq) phenolphthalein indicates OH- (pink) problem: why is ammonia (NH3) basic?

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) H+ is really H3O+ because water bonds with H+ hydronium ion H3O+ + = + base conj. acid 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− ] x [H3O+] = 10-14 [OH−] and [H3O+] are inversely proportional neutral water: Kw = [10-7] x [10-7] = 10-14 acidic: Kw = [10-9] x [10-5]= 10-14 basic: Kw = [10-3] x [10-11]= 10-14

[H3O+] ACIDS [H3O+] > 10-7 M 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 NaOH(s) → Na+(aq) + OH−(aq) OH− reduces 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 how acid precipitation forms. Understand the effects of acid precipitation and how they can be reduced. Understand how acid-base indicators work.

Acid Rain, Acid Fog acid rain: precipitation with a low pH (< 5) burning “high-sulfur” coal produce SO2 and SO3 that react w/ H2O to make H2SO3 and H2SO4 cars make NOX: reacts w/ H2O to make HNO2 and HNO3 corrodes metal dangerous to organisms decomposes limestone

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

Acid-Base Indicators ↔ compounds that respond to pH change by changing color contain a “weak acid” in an equilibrium 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 universal indicator: mixture, wide pH range

Plant Dyes and pH serviceberry, willow bark, Oregon grape root, contain indicators. These plants have traditionally been used as natural dyes for skins, feathers, and so on.

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 HNO3 (g) + H2O (l) → H3O+(aq) + NO3−(aq) weak acid: partially ionizes in water, reactants favored HC2H3O2(l) + H2O (l) ↔ H3O+(aq) + C2H3O2−(aq)

Acid Dissociation Constant (KA) HA ↔ H+ + A− Strong acids—high KA ( > 1, products favored) Weak acids—low KA ( < 1, reactants favored) Note that [H+] = [A− ] * use [H+] = 10-pH [H+] = [H3O+] [HA] = initial molarity – [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 [H+] (same value as [A-] ) [H+] = 10-pH = 10-1.93 = 0.012 M Determine [HA] [HA] = initial – [H+] = 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 (M) strong = product favored weak = reactant favored concentrated = lots of solute dilute = not much solute

Neutralization acid + base → salt + water H+ + OH− → H2O 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