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Acids and Bases

Acids and Bases. All you ever wanted to know, and more!. Definitions. Arrhenius acid – contains H + and ionizes in water. HCl + H 2 O  H 3 O + + Cl - hydronium ion - H + ion attached to water H 2 O + H + ↔ H 3 O +. Properties of Acids. Taste sour Electrolytes

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Acids and Bases

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  1. Acids and Bases All you ever wanted to know, and more!

  2. Definitions • Arrhenius acid – contains H+ and ionizes in water. HCl + H2O  H3O+ + Cl- hydronium ion - H+ ion attached to water H2O+ H+↔ H3O+

  3. Properties of Acids • Taste sour • Electrolytes • React with bases to form salt & water ie – HCl + NaOH  • React with active metals to produce hydrogen gas ie – HBr + Na 

  4. Polyprotic acids • Polyprotic - Acids that have many protons to donate i.e. H3PO4 • Monoprotic – one proton like HCl • Diprotic – two protons, like H2SO4 • H’s come off one at a time and require separate reactions for each H.

  5. example • H2SO4 + H2O  HSO4- + H+ + H2O • HSO4- + H2O  SO4-2 + H+ + H2O

  6. Review on Naming Acids • Binary Acids – • (2 elements, H and some nonmetal) • Hydro----ic • Ternary Acids – A. (3 elements, H & polyatomic ion) • “ate” ion - ___ic acid • “ite” ion - ___ous acid

  7. Arrhenius base – contains OH- and ionizes in water. NaOH + H2O  Na+ + OH- + H2O

  8. Properties of Bases • Taste bitter • Feels slippery • Reacts with acid to form salt and water • Electrolytes

  9. Brǿnsted-Lowry model • Acid – proton donor • Base – proton acceptor • What proton? H+ ion, once the electron is removed

  10. Conjugate acid-base pairs • Two substances related to each other by donating and accepting a single proton (H+) • Equilibrium reactions – reactions where the forward and reverse reactions can both occur. (between weak acids and bases)

  11. example HF + H2O ↔ H3O+ + F- acid base conjugate conjugate acid base NH3 + H2O ↔ NH4+ + OH- base acid conjugate conjugate acid base

  12. Water – acid, base or neutral? • Pure water is neutral because [H+] = [OH-] • Water can act like an acid or base depending on what it’s mixed with • Substances that can behave as both acid and base are said to be amphoteric.

  13. Auto Ionization of Water In pure water, 2 out of every billion molecules ionize according to this reaction: HOH  H+ + OH- This yields a concentration in water of: [H+] = 1 x 10-7 M [OH-] = 1 x 10-7 M

  14. So, all aqueous solutions (water) have both H+ and OH- ions present and the product, 1 x 10-14, is a constant. If you know [ ] of one you can calculate the other! [H+] [OH-] = 1 x 10 -14 constant

  15. Ion Concentrations [H+] [OH-] = 1 x 10-14 • What are the ion concentrations of a 0.000453M solution of HCl? • What are the ion concentrations of a 0.00250M solution of KOH?

  16. If…. [H+] > [OH-], solution is acidic [H+] = [OH-], solution is neutral [H+] < [OH-], solution is basic

  17. pH Scale • pH stands for power of Hydrogen • The pH scale was developed as an easier method of expressing ion concentrations. pH = -log[H+]

  18. pH Scale 0 acid 7 base 14 neutral

  19. [H3O+] [OH-] = 1 x 10-14 -log [OH-] -log [H3O+] antilog (-pOH) antilog (-pH) pH + pOH = 14

  20. Strengths of Acids and Bases • Strength of acids and bases is determined by how much they ionize (how much H+ or OH- they produce.) • Examples – HCl  H+ + Cl- ≈100% Therefore, HCl is considered strong.

  21. H2S ↔H+ + HS- >90% <10% H2S is considered a weak acid because not much H+ is produced

  22. So, how do you know if an acid or base is strong or weak? Most are weak, so memorize the strong acids and bases then assume everything else to be weak.

  23. Strong Acids • HCl, hydrochloric acid • HBr, hydrobromic acid • HI, hydroiodic acid • H2SO4,sulfuric acid • HClO4,perchloric acid • HClO3, chloric acid • HNO3, nitric acid

  24. Strong Bases • NaOH, sodium hydroxide • KOH, potassium hydroxide • RbOH, rubidium hydroxide • CsOH, cesium hydroxide • Ca(OH)2, calcium hydroxide • Sr(OH)2, strontium hydroxide • Ba(OH)2, barium hydroxide

  25. Once we know that an acid or base is weak, then what? • Weak acids and bases produce a solution containing a mixture of molecules and ions. The concentration of the ions is determined by using an equilibriumexpression.

  26. Equilibrium Expressions k = [products] [reactants] (except pure solids and liquids) Where: k = ionization constant [ ] = concentration in Molarity

  27. example HC2H3O2 + H2O(l ) ↔ C2H3O2- + H3O+ K = [C2H3O2-] [H3O+] [HC2H3O2]

  28. R I C E R = REACTION I = INITIAL concentration C = CHANGE in concentration E = EQUILIBRIUM concentration

  29. Example… 1. What are the ion concentrations of 0.5M HIO, ka = 2.3 x 10-11 2. What is the pH?

  30. Equilibrium always favors the weaker acid/base pair.

  31. Neutralization • When a strong acid reacts with a strong base to form a salt and water. • Example: 2NaOH + H2SO4 Na2SO4 + 2H2O

  32. Titration Problems • Titration – using a solution of known concentration to determine the concentration of an unknown solution • Standard solution – solution of known concentration

  33. Equivalence Point – the point when there are equal molar amounts of acid and base Indicator – substance that changes color as the pH changes. Indicators are chosen to change at the equivalence point, called the end point.

  34. example • In a titration, 42.8 mL of a standard solution of Ca(OH)2 is added to 20.5 mL sample of HCl. The concentration of the calcium hydroxide is 0.35 M. What is the molarity of the acid solution?

  35. Steps for Solving Problems MaVaH’s = MbVbOH’s

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