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

2. Definitions Arrhenius acid – contains H + and ionizes in water. 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 +

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

4. Polyprotic acids Polyprotic - Acids that have many protons to donate Polyprotic - Acids that have many protons to donate i.e. H 3 PO 4 Monoprotic – one proton like HCl Monoprotic – one proton like HCl Diprotic – two protons, like H 2 SO 4 Diprotic – two protons, like H 2 SO 4 H’s come off one at a time and require separate reactions for each H. H’s come off one at a time and require separate reactions for each H.

5. example H 2 SO 4 + H 2 O  HSO 4 - + H + + H 2 O H 2 SO 4 + H 2 O  HSO 4 - + H + + H 2 O HSO 4 - + H 2 O  SO 4 -2 + H + + H 2 O HSO 4 - + H 2 O  SO 4 -2 + H + + H 2 O

6. Review on Naming Acids 1. Binary Acids – A.(2 elements, H and some nonmetal) B.Hydro----ic 2. Ternary Acids – A. (3 elements, H & polyatomic ion) 1.“ate” ion - ___ic acid 2.“ite” ion - ___ous acid

7. Arrhenius base – contains OH - and ionizes in water. Arrhenius base – contains OH - and ionizes in water. NaOH + H 2 O  Na + + OH - + H 2 O

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

9. Brǿnsted-Lowry model Acid – proton donor Acid – proton donor Base – proton acceptor Base – proton acceptor What proton? 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 + ) 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) Equilibrium reactions – reactions where the forward and reverse reactions can both occur. (between weak acids and bases)

11. example HF + H 2 O ↔ H 3 O + + F - acid base conjugate conjugate acid base conjugate conjugate acid base acid base NH 3 + H 2 O ↔ NH 4 + + OH - NH 3 + H 2 O ↔ NH 4 + + OH - base acid conjugate conjugate base acid conjugate conjugate acid base acid base

12. Water – acid, base or neutral? Pure water is neutral because Pure water is neutral because [H + ] = [OH - ] Water can act like an acid or base depending on what it’s mixed with 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. 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 1. What are the ion concentrations of a 0.000453M solution of HCl? 2. 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 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 0 acid 7 base 14neutral

19. [H 3 O + ] [OH - ] = 1 x 10 -14 -log [H 3 O + ] pH + pOH = 14 antilog (-pOH) antilog (-pH) -log [OH - ]

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

21. 2 H + + HS - H 2 S ↔ H + + HS - > 90% <10% H 2 S 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 HCl, hydrochloric acid HBr, hydrobromic acid HBr, hydrobromic acid HI, hydroiodic acid HI, hydroiodic acid H 2 SO 4, sulfuric acid H 2 SO 4, sulfuric acid HClO 4, perchloric acid HClO 4, perchloric acid HClO 3, chloric acid HClO 3, chloric acid HNO 3, nitric acid HNO 3, nitric acid

24. Strong Bases NaOH, sodium hydroxide NaOH, sodium hydroxide KOH, potassium hydroxide KOH, potassium hydroxide RbOH, rubidium hydroxide RbOH, rubidium hydroxide CsOH, cesium hydroxide CsOH, cesium hydroxide Ca(OH) 2, calcium hydroxide Ca(OH) 2, calcium hydroxide Sr(OH) 2, strontium hydroxide Sr(OH) 2, strontium hydroxide Ba(OH) 2, barium 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 equilibrium expression. Weak acids and bases produce a solution containing a mixture of molecules and ions. The concentration of the ions is determined by using an equilibrium expression.

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

27. example HC 2 H 3 O 2 + H 2 O( l ) ↔ C 2 H 3 O 2 - + H 3 O + K = [C 2 H 3 O 2 - ] [H 3 O + ] [HC 2 H 3 O 2 ]

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, k a = 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. When a strong acid reacts with a strong base to form a salt and water. Example: Example: 2NaOH + H 2 SO 4  Na 2 SO 4 + 2H 2 O

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

33. Equivalence Point – the point when there are equal molar amounts of acid and base 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. 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? 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 M a V a H’s = M b V b OH’s