1. Chapter 4 Notes Types of Chemical Reactions and Solution Stoichiometry

3. Things about water that are neato.. high specific heat high heat of vaporization high cohesive and adhesive forces Solid water floats – think about everything else that doesn’t!

4. Why is water so cool? Its structure. The electrons are more attracted to the oxygens than the hydrogens (the oxygen is more electronegative) and so it creates a partial charge. This partial charge makes water extremely polar. Electron space hogs on the oxygen.

5. Why a lot of ionic things dissolve in water and why some don’t. Opposite charges of ions that make up ionic compounds are attracted to the positive and negative portions of a water molecule. However, when the attraction of the ions to each other is greater than that to the water, the substance is insoluble.

6. What about nonionic substances? Many nonionic substances can dissolve as well. If a molecule contains a polar bond, then the substance should be at least moderately soluble. Examples: alcohols, sugar Nonionic substances that don’t dissolve are non-polar or have extremely large non-polar areas. Examples: fats, alkanes, long alcohols Remember that “like dissolves like”

7. Electrolytes An electrolytic solution is an aqueous solution that is able to conduct electricity because of the presence of ions. Strong electrolytes – are solutions made of compounds that completely dissociate. Examples: soluble salts, strong acids and bases Weak electrolytes – are solutions made of compounds that slightly dissociate. Examples: solutions with very few ions, weak acids and bases

8. Nonelectrolytes Substances that doesn’t conduct electricity. The solute may dissolve, but the dissolving does not produce ions. Examples: sugar, alcohol, starch What happens eletrolytically to a barium hydroxide solution as it is slowly titrated with sulfuric acid?

9. Composition of Solutions Molarity – concentration unit in moles of solute per liter of solution.

10. Exercise 1 Calculate the molarity of a solution prepared by dissolving 11.5 g of solid NaOH in enough water to make 1.50 L of solution.

11. Exercise 2 Calculate the molarity of a solution prepared by dissolving 1.56 g or gaseous HCl in enough water to make 26.8 mL of solution.

12. Exercise 3 Give the concentration of each type of ion in the following solutions a. 0.50 M Co(NO 3 ) 2 b. 1 M Fe(ClO 4 ) 3

13. Exercise 4 Calculate the number of moles of Cl - ions in 1.75 L of 1.0 x 10 -3 M ZnCl 2.

14. Exercise 5 Typical blood serum is about 0.14 M NaCl. What volume of blood contains 1.0 mg NaCl?

15. Exercise 6 To analyze the alcohol content of a certain wine, a chemist needs 1.00 L of an aqueous 0.200 M K 2 Cr 2 O 7 (potassium dichromate) solution. How much solid K 2 Cr 2 O 7 must be weighed out to make this solution?

16. Making a molar solution (a) A weighed amount of a substance (the solute) is put into the volumetric flask, and a small quantity of distilled water is added. (b) The solid is dissolved in the water by gently swirling the flask (with the stopper in place). (c) More water is added, until the level of the solution just reaches the mark etched on the neck of the flask (d). If you dump solid into the total volume of water you are neglecting the space the solid will occupy and your solution will NOT be correct.

17. Dilutions A more concentrated solution is already made. Use M 1 V 1 =M 2 V 2 Measure volume with a pipet—a device for accurately measuring and transferring a given volume of solution graduated pipet (a)—used to measure volumes for which a volumetric pipet is not available— has lines. volumetric or transfer pipet (b)—gives only one measurement 5mL, 10mL, etc.

18. Making a molar solution (part II) (a) A volume of a substance (the solute) is put into the volumetric flask using a pipet, and a small quantity of water is added. (b) If a strong acid or base is being diluted, place some distilled water in the flask first before pipetting. (c) Add distilled water until the level of the solution just reaches the mark etched on the neck of the flask (d). Adding two measures of water may NOT give the total water. Ex. 30mL + 30mL ≠ 60mL.

19. Exercise 7 What volume of 16 M sulfuric acid must be used to prepare 1.5 L of a 0.10 M H 2 SO 4 solution ?

20. Types of Chemical Reactions Balancing – do I really need to explain this? Why? Conservation of matter Describing Reactions complete balanced equation—gives the overall reaction stoichiometry, but NOT the forms of the reactants & products as they exist in solution. complete ionic equation—represents as IONS all reactants & products that are strong electrolytes net ionic equations—includes only those solution components undergoing a change. Spectator ions are NOT included. This is what they usually ask for on the AP test. spectator ions--not involved in the reaction process. They started as an ion AND finished as an ion.

21. Solubility Rules – KNOW THESE! The following compounds are soluble: All salts of alkali metals (Na,K,…) and NH 4 + Hydroxides of Groups 1A and 2A Except Be 2+ Mg 2+ All Salts of Cl - Br - I - Except Ag + Hg + Pb 2+ Salts of F - Except Mg 2+ Ca 2+ Sr 2+ Ba 2+ Pb 2+ Salts of Sulfates (SO 4 2- ) Except Ag + Hg + Pb 2+ Sr 2+ Ba 2+ Almost all salts of Nitrate (NO 3 - )Chlorate (ClO 3 - ) Perchlorate (ClO 4 - ) Acetate (C 2 H 3 O 2 - )

22. And this too! The following compounds are generally not soluble: Carbonates (CO 3 2- )Phosphates (PO 4 3- ) Oxalates (C 2 O 4 2- )Chromates (CrO 4 2- ) Sulfides (S 2- )Oxides (O 2- )

23. Molecular Snap-shot

24. Exercise 8 Using the solubility rules, predict what will happen when the following pairs of solutions are mixed. a. KNO 3 (aq) and BaCl 2 (aq) b. Na 2 SO 4 (aq) and Pb(NO 3 ) 2 (aq) c. KOH (aq) and Fe(NO 3 ) 3 (aq)

25. Exercise 9 For each of the following reactions, write the molecular equation, the complete ionic equation, and the net ionic equation. a. Aqueous potassium chloride is added to aqueous silver nitrate to form a silver chloride precipitate plus aqueous potassium nitrate.

26. Exercise 9 - continued b. Aqueous potassium hydroxide is mixed with aqueous iron(III) nitrate to form a precipitate of iron(III) hydroxide and aqueous potassium nitrate.

27. Exercise 10 - stoichiometry Calculate the mass of solid NaCl that must be added to 1.50 L of a 0.100 M AgNO 3 solution to precipitate all the Ag + ions in the form of AgCl.

28. Exercise 11 - stoichiometry When aqueous solutions of Na 2 SO 4 and Pb(NO 3 ) 2 are mixed, PbSO 4 precipitates. Calculate the mass of PbSO 4 formed when 1.25 L of 0.0500 M Pb(NO 3 ) 2 and 2.00 L of 0.0250 M Na 2 SO 4 are mixed.

29. Acid-Bases Reactions Acids – any compounds that on reaction with water produces a hydronium ion (hydrogen ion) Base – any compound that provides a hydroxide (there are exceptions) Neutralization – when stochiometrically equal amounts of acid and base are added together. The resulting solution is not always neutral. (more later)

30. Exercise 12 – acid-base What volume of a 0.100 M HCl solution is needed to neutralize 25.0 mL of 0.350 M NaOH ?

31. Exercise 13 – acid-base In a certain experiment, 28.0 mL of 0.250 M HNO 3 and 53.0 mL of 0.320 M KOH are mixed. Calculate the amount of water formed in the resulting reaction. What is the concentration of H + or OH - ions in excess after the reaction goes to completion?

32. Analysis Gravimetric Analysis – determining the amount of a substance by precipitation reactions. Volumetric Analysis - determining the amount of a substance by titrating. Colorimetric Analysis - determining the amount of a substance by analysis with a spectrophotometer.

33. Titrating Vocab. Review titrant—the substance delivered from a buret so that its volume is accurately known analyte—the substance being analyzed; its mass or volume must also be accurately known equivalence point--# moles of OH - equals (is equivalent to) # moles of H 3 O + [acid-base titration; redox titrations also exist!] indicator--undergoes a color change near the equivalence point. standardization--a procedure for establishing the exact concentration of a reagent.

34. Exercise 14 – titration A student carries out an experiment to standardize (determine the exact concentration of) a sodium hydroxide solution. To do this the student weighs out a 1.3009-g sample of potassium hydrogen phthalate (KHC 8 H 4 O 4, often abbreviated KHP). KHP (molar mass 204.22 g/mol) has one acidic hydrogen. The student dissolves the KHP in distilled water, adds phenolphthalein as an indicator, and titrates the resulting solution with the sodium hydroxide solution to the phenolphthalein endpoint. The difference between the final and initial buret readings indicates that 41.20 mL of the sodium hydroxide solution is required to react exactly with the 1.3009 g KHP. Calculate the concentration of the sodium hydroxide solution.

35. Exercise 15 – titration An environmental chemist analyzed the effluent (the released waste material) from an industrial process known to produce the compounds carbon tetrachloride (CC1 4 ) and benzoic acid (HC 7 H 5 O 2 ), a weak acid that has one acidic hydrogen atom per molecule. A sample of this effluent weighing 0.3518 g was shaken with water, and the resulting aqueous solution required 10.59 mL of 0.1546 M NaOH for neutralization. Calculate the mass percent of HC 7 H 5 O 2 in the original sample.