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Acids / Bases

Acids / Bases. Acids Taste sour Corrosive Conducts electricity pH < 7 Turns litmus paper red Formula contains H- Produces hydronium ions(H 3 O + ) in solution. Bases Taste Bitter Slippery to the touch Corrosive Conducts electricity pH > 7 Turns litmus paper blue

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Acids / Bases

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  1. Acids/Bases

  2. Acids Taste sour Corrosive Conducts electricity pH < 7 Turns litmus paper red Formula contains H- Produces hydronium ions(H3O+) in solution Bases Taste Bitter Slippery to the touch Corrosive Conducts electricity pH > 7 Turns litmus paper blue Formula contains –OH Produces hydroxide ions(OH-) in solution Turns phenolphthalein pink

  3. pH Scale BASES ACIDS 14 0 Strong NEUTRAL Strong 7 IONS – More ions = stronger A/B

  4. Strong Acids • Completely ionize when put in water H2SO4 HClO4 HNO3 HCl HBr HI Strong Bases • Completely ionize when put in water NaOH KOH RbOH CsOH Ca(OH)2 Sr(OH)2 Ba(OH)2 Any other acid/base not listed here is Weak Weak Acids/Bases – partly ionize in water

  5. Naming Acids/Bases BASES Name of metal + name of –OH EX: NaOH sodium hydroxide ACIDS Binary 2 elements H + nonmetal Prefix – hydroRoot – acid name Suffix – ic EX: HCl Hydrochloric acid Oxyacid H +polyatomic ion PI ends –ate=ic PI ends –ite = ous EX: H2SO4 sulfuric acid H2SO3 sulfurous acid

  6. Acid/Base Theories Arrhenius – in water only Acids – produce H+(H3O+)Bases – produce OH- Bronsted-Lowry Acids – proton (H+) donorBases – proton acceptor Lewis – broadest view Acids – electron pair acceptorBases – electron pair donor

  7. H+ + NH3 NH4+ 2. NaOH  Na+ + OH- 3. HF + H2O  F- + H3O+ • NH3 + H2O  OH- + NH4+ • HCl  H+ + Cl-

  8. Bronsted-Lowry • Acid – what it makes is its conjugate base Strong acid = weak conjugate base (and vice versa) • Base – what it makes is its conjugate acid Strong base = weak conjugate acid (and vice versa) NH3 + H2O  OH- + NH4+

  9. Practice • NH4+ + OH- NH3 + H2O • HBr + H2O  H3O+ + Br- • CO3-2 + H2O  HCO3- + OH- • HSO4- + H2O  H3O+ + SO4-2

  10. Neutralization Acid + Base  Salt+ water HCl + Sr(OH)2 SrCl2 + H2O **Double displacement reaction (Ionic Compound) (HOH)

  11. Self Ionization of Water H2O + H2O  H3O+ + OH- In pure water, H3O+ OH- In a solution, H3O+ OH- When H3O+> OH-, the solution is acidic When H3O+<OH-, the solution is basic

  12. Kw = [H3O+][OH-] = 1.0 x 10-14 pH = -log [H3O+] ; pOH = -log [OH-] pH + pOH = 14 [] = concentration = Molarity If a solution contains 1.3 x 10-2 M OH-, what is the hydronium ion concentration? What is the pH of the solution? Using the above equations: 1.0 x 10-14 = [x][1.3 x 10-2] [x] = [H3O+] = 7.69 x 10-13M pH = -log [7.69 x 10-13] pH = 12.1 pOH = -log [1.3 x 10-2] pOH = 1.89 pH + 1.89 = 14 pH = 12.1 OR

  13. Determine the [H3O+], [OH-], and the pH of a 0.01 M Sr(OH)2 solution. **Sr(OH)2 is a strong base, so it completely breaks apart into its ions when put in water – write the equation Sr(OH)2 (s)  Sr+2(aq) + 2OH-1(aq) Sr(OH)2 produces OH- when put into solution and since we know the [] of Sr(OH)2, we can use the mole to mole ratio to determine the concentration of the OH- Sr(OH)2(s)  Sr+2(aq) + 2OH-1(aq) 0.01 M 0.02 M [OH-] = 0.02 M Use Kw to find [H3O+] 1.0 x 10-14 = [H3O+][0.02] [H3O+] = 5.0 x 10-12M pH = -log (5.0 x 10-12) pH = 11.3

  14. If a solution has a pH of 4.23, what is the concentration of [H3O+]? What will the [OH-] be?

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