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Acids and Bases
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  1. Acids and Bases L

  2. Properties of acids • Sour • Conduct electricity • Strong electrolytes (some) • Weak electrolytes (some) • React with metals to form hydrogen gas • React with hydroxides to form water and salts

  3. Properties of bases • Bitter • Feel slippery • Can be strong or weak electrolytes • React with acids to form water and salts • Change indicators (litmus turns blue)

  4. Naming Acids (H) Bases • If the  ends in “-ide” the name is “Hydro-,” the root of , the suffix “-ic” then “acid” Ex: HCl = Hydro-chlor-ic acid HF = Hydro-flor-ic acid • If  ends in “-ate” the name is the root of , and the suffix “-ic” Ex: H2SO4 = sulferic acid • If  ends in “-ite” the name is the root of , and the suffix “-ous” Ex: HNO2 = Nitrous acid (SEE NEXT) • Compound that forms “–OH” in water. • Naming ionic compounds

  5. Lets Practice!! • Name an acid or base when given the formula and vice versa. A. H2SO4 E. Magnesium Hydroxide B. HCl F. Phosphoric Acid C. NaOH G. Sulfurous Acid D. HNO3 H. Hydrobromic Acid

  6. Acid Base theories • Arrhenius • Acid – creates H+/H3O+ as the only positive ion • HCl(aq) H+(aq) + Cl-(aq) • Base – creates OH- as the only negative ion • NaOH(aq) Na+(aq) + OH-(aq) • Neutralization – mixing of an acid with a base that produces salt and water HCl(aq) + NaOH(aq) H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq)  NaCl(aq) + H2O(l) H+(aq) + OH-(aq)  H2O(l)

  7. Acid Base Theories • Arrhenius: Limitations NH3(g) + HCl(g) NH4+ + Cl-NH4Cl(s) • Since this reaction doesn’t produce hydronium or hydroxide what is it??? • Bronsted- Lowery: • Acid – proton donator • Base – proton acceptor • Coordinate covalent bond is formed • One atom provides both electrons for a covalent bond (the base)

  8. The great proton grab • Bronsted-Lowery • Acids and bases form conjugate acid-base pairs • Whichever base is stronger will exist in higher concentrations at equilibrium.

  9. Water • Amphiprotic – can be either an acid or base • Self ionization • Water falls apart into ions • H2O H+ + OH- • Very small amount • [H+ ] = [OH-] = 1 x 10-7M • A neutral solution. • In water Kw = [H+ ] [OH-] = 1 x 10-14 • Kw is called the water dissociation constant. • In water, If you know either [H] or [OH] you can calculate the other

  10. Ion Product constant • H2O H+ + OH- • Kw is constant in every aqueous solution [H+] x [OH-] = 1 x 10-14M2 • If [H+] > 10-7 then [OH-] < 10-7 • If [H+] < 10-7 then [OH-] > 10-7 • If we know one, we can determine the other. • If [H+] > 10-7acidic [OH-] < 10-7 • If [H+] < 10-7basic [OH-] > 10-7

  11. Examples: • Is The solution Acidic, Basic or Neutral? What is the [OH-]/ [H+]? 1. [H+] = 1.0 X 10-5 2. [H+] = 1.0 X 10-10 3. [OH-] = 1.0 X 10-2 4. [H+] = 1.0 X 10-7 5. [OH-] = 1.0 X 10-7

  12. Logarithms • Powers of ten. • A shorthand for big or small numbers. • pH = -log[H+] • in neutral pH = - log(1 x 10-7) = 7 • in acidic solution [H+] > 10-7, 100→10-6.99 • pH < -log(10-7) • pH < 7 • in base pH > 7 • pH > -log(10-7) • [H+] < 10-7, 10-7.01→10-14

  13. pH and pOH • pOH = - log [OH-] • [H+] x [OH-] = 1 x 10-14M2 • pH + pOH= 14

  14. 100 10-1 10-3 10-5 10-7 10-9 10-11 10-13 10-14 14 13 11 9 7 5 3 1 0 10-14 10-13 10-11 10-9 Basic 10-7 10-5 10-3 10-1 100 [H+] pH 0 1 3 5 7 9 11 13 14 Acidic Neutral Basic pOH [OH-]

  15. Examples: • Calculating pH from [H+] [H+] = 1.0 X 10-6 mol/L [H+] = .0001M [OH-] = 1.0 X 10-2 mol/L [OH-] = 1.0 X 10-11 mol/L • Using pH to find [H+] • A. 4.0 • B. 11.0 • C. 8.0

  16. Strength vs Concentration • Strong: dissociates fully in water • Example: HCl, HBr, HClO4, HNO3, H2SO4 • For every mole of acid placed in solution you get a mole of H+ • Monoprotic – can donates only 1 proton • polyprotic – can donate more then 1 proton • Weak: dissociated partially in water • Established equilibrium (Ka or Kb) • Donates one proton but can usually give more

  17. Using Ka/b • Ka – acid ionization constant • If you know the initial concentration of an acid and the formula you can calculate [H+] given Ka • Same equation as Keq • Ex: CH3COOH + H2O  CH3OOH- + H3O+ What is [H3O+], if 2.0M of acetic acid was dissolved in water and has a Ka of 1.8 x 10-5?

  18. CH3COOH + H2O  CH3OOH- + H3O+What is [H3O+], if 2.0M of acetic acid was dissolved in water and has a Ka of 1.8 x 10-5? • Ka = [CH3OOH-][H3O+] [CH3COOH] • Ka = 1.8 x 10-5 = [CH3OOH-][H3O+] [CH3COOH] 1.8 x 10-5 = x(x) 2M 1.8 x 10-5 = X2 2M

  19. CH3COOH + H2O  CH3OOH- + H3O+What is [H3O+], if 2.0M of acetic acid was dissolved in water and has a Ka of 1.8 x 10-5? 1.8 x 10-5 = X2 2M 3.6 x 10-5 = X2 √(3.6 x 10-5) = √(x2) = [H3O+] 6.0 x 10-3 = [H3O+] NOTE: if an acid has a Ka it is WEAK! Strong Acids: HCl, HBr, HI, HNO3, HClO3, HClO4, H2SO4

  20. Indicators – dyes that change color in the presence of an acid or base Hydrangea Red Cabbage • If grown in acidic soil turns pink • If grown in basic soil turns blue • Chopped and boiled, remaining liquid • Turns pink in acids • Turns green in bases

  21. Indicators (common) Litmus Paper Phenolphthalein • Red Litmus paper • Tests for bases • Acid – no change • Base – turns blue • Blue Litmus paper • Tests for acids • Acid – turns red • Base – no change • Neutral Litmus paper • Tests for both • Acid – turns red • Base – turns blue • Active ingredient in laxatives • Acid – Clear and colorless • Bases – Pink • Used in titrations

  22. Titrations • Procedure for calculating the pH of an acid/base by reacting it with a known concentration of base/acid • (Macid)(Vacid)(# of H+) = (Mbase)(Vbase)(# of OH-)

  23. Buffers or buffer solutions • Resist change in pH caused by the addition of acids or bases • Contains something that reacts with both acids and bases • Mixtures of weak acids and bases • Conjugate pairs • Most common in organic systems • H2CO3/HCO3- • Non-conjugate acid-base pairs • Most common in pharmaceutical (Alka-seltzer) • NH4+/CH3COO- • Amphoteric species • Water • HCO3-, HPO4-2 in blood pH

  24. Oxidation and Reduction reactionsLEO goes GER Oxidation Reduction • the loss of electrons from atoms of a substance • LEO- Loss of Electrons is Oxidation • Na → Na+ + e- • Na loses an electron and thus, is oxidized. • Oxidizing agent- substance that facilitates oxidation by accepting lost electrons; it is the substance that is reduced. • the gain of electrons by atoms of a substance • GER- Gain of Electrons is Reduction • Cl2 + 2e- → 2Cl- • Cl2 gains electrons and thus, is reduced. • Reducing agent- the substance that facilitates reduction by losing electrons; it is the substance that is oxidized.

  25. 2KBr + Cl2 → 2KCl + Br2 • Net ionic equation: 2Br - + Cl2 → Br2 + 2Cl – • Half reaction: 2Br - → Br2 + 2e – • 2Br– is oxidized to become Br2. • 2Br– is the reducing agent • Half reaction: Cl2 + 2e - → 2Cl – • Cl2 is reduced to become 2Cl – • Cl2 is the oxidizing agent