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Acids, Bases, and Salts

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. ACIDS pH < 7 sour taste electrolytes react with metals to make hydrogen gas

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Acids, Bases, and Salts

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  1. Acids, Bases, and Salts

  2. 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.

  3. 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 Acids

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

  5. Arrhenius Acids Arrhenius discovered that acids are: • molecules that contain H • ionize in water to make H+ • HCl→ H+(aq) + Cl–(aq) • H2SO4→ H+(aq) + HSO4–(aq) • HBr → ? • H3PO4 → ? Svante Arrhenius 1859-1927

  6. 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.

  7. Acid Nomenclature • USE YOUR YELLOW SHEET! • use the stem and ending of the anion name -ide hydro-stem-ic acid -ate stem-ic acid -itestem-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)

  8. 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) Ca(OH)2(s) → ? KOH(s) → ? • phenolphthalein indicates OH- (pink) • problem: why is ammonia (NH3) basic?

  9. 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

  10. 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

  11. 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.

  12. 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

  13. [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

  14. 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

  15. Objectives • Understand how acid precipitation forms. • Understand the effects of acid precipitation and how they can be reduced. • Understand how acid-base indicators work.

  16. Acid Rain, Acid Fog • acid rain/fog: 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

  17. Acid Rain in the USA

  18. 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

  19. Acid-Base Indicators • compounds that respond to pH change by changing color • contain a “weak acid” in a chemical 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

  20. Plant Dyes and pH • serviceberry, willow bark, Oregon grape root, are indicators • have been used as natural dyes for skins, feathers, etc.

  21. 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.

  22. 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)

  23. 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+]

  24. 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

  25. 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.

  26. 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

  27. 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)

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

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

  30. 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

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