1. Final Exam Review

2. Oxoacids Central element, surrounded by oxygens e.g. (not all listed) Recognizing Acids Hydrohalic Acids Group 7 ion with H+ HF, HCl, HBr, HI Carboxylic acids - COOH group

3. Recognizing Bases First or second groups combined with an OH group. Examples: NaOH KOH Mg(OH) 2 Ect…..

4. Conjugate Acid/Base pairs Acids lose a proton to become the conjugate base. Bases gain a proton to become it’s conjugate acid All acid/base reactions have both an acid and a base in them? So what is the base and conjugate acid in the first reaction? What is the acid and conjugate base in the second reaction? Acid Conjugate Base Base Conjugate Acid Conjugate Base

5. Acid/Base Conceptual Understanding Questions: In what range must the pH of a 0.05M solution of a weak acid fall? What must be true about the [H+] of a weak acid solution? What must be true about the pH? What must be true about the [OH-] of a weak acid solution? What must be true about the pH? Less than 7: Greater than: -log(0.05M)=1.3 greater than 10 -7 : Less than 7 Less than 10 -7 : Greater than 7

6. Example: Ranking without using Ka Rank the following solutions in order of increasing acidity. Assume the concentration is the same for all solutions and that all are low enough concentrations to be soluble. You do not need to look up anything to do this problem. HCl, NaOH, Ca(OH) 2, N(CH 3 ) 2 H, N(CH 3 ) 2 H 2 + Strong Acid Strong Base Strong Base x2 hydroxides Weak base Weak acid HCl Ca(OH) 2 NaOH N(CH 3 ) 2 HN(CH 3 ) 2 H 2 +

7. Re-deriving: Important Relations Can you just memorize this without understanding it? What happens if the temperature changes? pH=7 @25 o C ONLY Taking the log of both sides:

8. Kw Example Problem: Commonly missed on the midterm If water is heated and placed under pressure it has a pH of 6.8. Find the Kw of water under these conditions. Is it acidic, basic or neutral? Important hints: what is the definition of neutral (no, its not pH=7, that is only true at 25 C, what is the actual definition of neutral? ***study tip: what else do you know because of the pH being lower? What does this say about the thermodynamics? Kw? [OH-] and [H+] ect….?

9. Weak and Strong Acid and Bases: Compare and Contrast For strong acids and bases, we assume they completely ionize. Equilibrium lies far to the right Equilibrium lies far to the left Weak acids and bases ionize to a very limited extent. Very Large Ka Very small Ka How do I know which is strong and weak? Memorize the strong, the rest are weak

10. The larger the K a the stronger the acid. vs The larger the K b the stronger the base.

11. Percent Ionization: Pictorial Representation Original Acid: 14 Dissociated Acid: 3 % ionization= 3/14*100%= 21.4%

12. Percent Ionization: Example The percent dissociation of a 0.800 M aqueous monoprotic weak acid is 0.10%. What is the Ka value for the acid? ***study tip: make sure you could do this same problem if given the Ka and asked to find the percent dissociation (aka, do it backwards) ICEICE 0.800 0 0 0.08 +0.08 -0.08 0.08 0.8-.08

13. K w and the relationship between K a and K b To convert between K a and K b use K w As K a increases K b of it’s conjugate base decreases

14. Example: Find the Kb of HCO 3 - ?

15. Molecular Structure and Strength of Acids Based on Bond enthalpy weak strong Hydrohalic Acids Oxo and carboxylic acids:The more you stabilize the anion, the stronger the acid Oxoacids Strength increases with increasing electronegativity of central atom. Example: HClO 3 >HBrO 3 Increases as oxidation number of central atom increases. Example: HClO 4 >HClO 3 >HClO 2 >HClO Carboxylic acids (Electronegative atoms in R group) > >

16. Buffers: How to make them Acid and its conjugate base, or a base and its conjugate acid Essentially the systems we’ve talked about in the common ion effect questions Examples: approximately equimolar mixtures of HCN and NaCN, HF and NaF, NH 4 Cl and NH 3. Pick a buffer where the pKa is near the pH of the buffer you desire. Make the conjugate acid base pair by reacting some of the material Examples: H 2 CO 3 with an appropriate amount of NaOH which will react to form the conjugate base. Must have some of each after reaction is complete: Example: 1 mol carbonic acid with 0.5 moles NaOH.

17. Buffers: Why do they work It works by converting a strong acid into a weak acid, or a strong base into a weak base. A strong base can’t exist in solution with a weak acid it must react A strong acid can’t exist in solution with a weak base it must react Examples: aka

18. A 100 mL buffer solution is 0.100M Nitrous acid and 0.100M Sodium nitrite. Calculate the pH if a)0.004 moles of NaOH is added to the solution b) calculate the pH of 0.0150 mols of NaOH is added to the solution. Assume no change in volume. IFIF 0.005 M0.010M - 0.004 M + 0.004 M 0.006 M 0. M + 0.004 M It’s a buffer! 0.0150 M0.010M - 0.010 M + 0.010 M 0.005 M 0. M+ 0.01 M It’s not a buffer! But we still need to find pH.

19. For each titration state what is the main species present at each step in the titration. Use HA as the weak acid and A- as the weak base. HA HA/A- A-

20. Titration Example Problem Suppose that 10.0g of an unknown monoprotic weak acid, HA, is dissolved in 100mL of water. To reach the equivalence point, 100.0mL of 0.10 M NaOH was used. After the addition of 50.0 mL, the pH of the solution was found to be 4.00. a)State how many moles of acid were added initially b)what is the molar mass of the acid? c)what is the value of pKa for the acid? d)What is the pH at the equivalence point(assume that auto ionization of water assumption is valid)?

21. Molar Solubility- The amount of a compound that is soluble. Mols/L Calculated from K sp Usually designated “s”, I use x as we’ve been doing with equilibrium constants

22. Common Ion Effect and Solubility Just like in K eq, if you add it to a salt solution with a common ion it changes the solubility Fill in concentration to K sp equation You may make an ice chart if it helps, but it is not necessary.

23. Common Ion Effect Conceptual Questions: What would each of the solutions do to the solubility of Calcium hydroxide as compared to pure water? Explain each. A 0.1 M solution HCl Increase: H+ removes OH- from solution. Le Chatliers shifts to right A 0.1M solution NaOH Decrease: Increased OH- shifts reaction to the left. A 0.1M solution calcium nitrate. Decrease: Increased calcium shifts reaction to the left. A solution buffered at pH 10. Decrease: pH 10=basic= higher OH- concentration, shifts reaction left. A solution of NaCl No Effect: No common ion in NaCl. Applicable solubility rules: Ca(OH) 2 is insoluble, Nitrate salts are soluble.

24. Common Ion Effect and Solubility You can have two different salt solutions that are perfectly soluble, but when mixed form a precipitate You can calculate if this will happen Q<K no precipitate Q=K saturated solution Q>K precipitate We can see if you mix two salt solutions if a precipitate will form. This is like using Q to determine which way the reaction will go.

25. Example: Common Ion Effect and Solubility Q<K no precipitate Q=K saturated solution Q>K precipitate A 10mL solution of 0.100 M sodium bromide and a 10mL solution of 0.1M silver nitrate are mixed. Does a precipitate form? (no calculation needed) A 10mL solution of 1.00x10 -5 M sodium bromide and a 10mL solution of 1.00x10 -5 silver nitrate are mixed. Does a precipitate form? A 10mL solution of 1.00x10 -7 M sodium bromide and a 10mL solution of 1.00x10 -7 silver nitrate are mixed. Does a precipitate form? Well above K so yes! Q>K so yes Q<K so no