1 Unit 4a: Solution Stoichiometry Last revised: October 19, 2011 If you are not part of the solution you are the precipitate. You should be able to: Vocabulary of water solubility Differentiate between strong and weak electrolytes Calculate molarity Calculate molarity of ions in solution Calculate using dilution equation Explain how to dilute a concentrated substance Use stoichiometry to convert to and from molarity Know basics of acids and bases Solve neutralization stoichiometry3 Explain what a titration is Solve titration math Assign oxidation numbers Balance redox reactions using the half reaction method Balance redox reactions in an acid Balance redox reactions in a base
2 4.1 Water, the common Solvent Read and outline 4.1 Define Polar Molecule Hydration Solubility Why is ethanol soluble in water? What is a useful rule to follow when predicting solubility? 4.1 Notes Water is a polar water molecule. Above you will see δ + or δ -. This symbol is used to denote a partial charge of the electronegativity within the atom. Electronegativity is when an atom pulls the electrons from a shared covalent bond closer to itself making it partially more negative. In this case, the oxygen pulls the electrons in and the hydrogen atoms become more positive. During hydration, the partially negative oxygen surrounds the positive salt ions (cations), and partially positive hydrogen atoms surround the negative salt ion (anion).
3 4.2 The Nature of Aqueous Solutions: Strong and Weak Electrolytes Read and outline 4.2 Define Solute Solvent Electrical conductivity Strong electrolyte Weak electrolyte Nonelectrolyte Strong acid Strong base Weak acid Weak base Explain the ionic theory in your own words. What are the 7 strong acids? How do you figure out if a base is strong or weak? What are the strong bases?
4 4.2 Notes Another word for a liquid homogenous mixture is solutions. A solution is formed from a solute and solvent. The solute is the substance being dissolved (salt, acid, base.), and the solvent is what is doing the dissolving (water for our purposes). One useful property for characterizing a solution is its electrical conductivity, its ability to produce an electric current. A strong electrolyte will create a current in the solution. When a light fixture is added to the solution a bright light emits. Strong electrolytes are soluble salts, strong acids, and strong bases. A weak electrolyte emits a dim light; weak acids and weak bases cause this. Nonelectrolytes emit no light because there are no ions to create a current in the solution. Soluble organic compounds and insoluble salts fall in this category. 4.3 The composition of Solutions Read and outline 4.3 Define Molarity Standard solution Dilution How do you make a standard solution? What is a stock solution? Why should you never ever ever ever ever ever pour liquid back into a stock solution? 4.3 Notes We will express concentration in terms of molarity in this class. Molarity is the amount a moles per 1 liter of solution. Equation: M = molarity, n = moles, V = Volume in liters
5 Example 4.3a Calculate the molarity of a solution prepared by dissolving 11.5 g of solid NaOH in enough water to make 1.50L of solution. Example 4.3b Give the concentration of each type of ion in the following solution: 0.50 M Co(NO 3 ) 2 and 1M Fe(ClO 4 ) 3. Example 4.3c Calculate the number of moles of Cl - ions in 1.75 L of 1.0 x 10-3 M ZnCl 2. Example 4.3d Typical blood serum is about 0.14 M NaCl. What volume of blood contains 1.0mg NaCl? Example 4.3e To analyze the alcohol content of certain wine, a chemist needs 1.00 L of an aqueous 0.200 M K 2 Cr 2 O 7 solution. How much solid potassium dichromate must be weighted out to make this solution?
6 To save time and space, solutions are prepared in concentrated from called a stock solution. Water is then added to achieve the desired molarity for a lab. This process is called dilution. A typical dilution calculation involves determining how much water must be added to an amount of stock solution. The key to doing these calculations is to remember that the number of moles after dilution is the equal to the number of moles before the dilution. The reason why is because only water is added to accomplish the dilution. A proper dilution procedure involves two types of glassware: a pipet, and a volumetric flask. You will first need to find the amount of concentrated solution added to the water. Equation: M = Molarity, V = Volume, 1 = first solution, 2 = second solution Example 4.3f How would you dilute 16M sulfuric acid to 1.5L of 0.10 M sulfuric acid solution? Example 4.3g Write a procedure on how to prepare the dilution from example 4.3f.
7 Worksheet 4b.01 Molarity Name 1. Calculate the molarity of each of the following solutions: a. A 16.45 g sample of NaCl is dissolved in enough water to make 1.000 L of solution b. A 853.0 mg sample of KIO 3 is dissolved in enough water to make a 250.0 ml solution. c. A 0.4508 g sample of iron is dissolved in a small amount to concentrated nitric acid forming Fe +3 in solution and is diluted to a total volume of 500.0 ml. Calculate the molarity of Iron (III). 2. Calculate the concentration of all ions present in each of the following solutions a. 0.100 mol of Ca(NO 3 ) 2 in 100.0 ml of solution b. 2.5 mol of sodium sulfate in 1.25 L of solution c. 5.00 g of ammonium chloride in 500.0 ml of solution d. 1.00 g of potassium phosphate in 250.0 ml of solution 3. A stock solution containing Mn +2 ions was prepared by dissolving 1.584 g pure manganese metal in nitric acid and diluting to a final volume of 1.000 L. The following solutions were prepared by diluting: For solution A, 50.00 ml of stock solution was diluted to 1000.0 ml. For solution B 10.00 ml of solution A was diluted to 250.0 ml. For solution C, 10.00 ml of solution b was diluted to 500.0 ml. Calculate the concentrations of the stock solution and solutions A, B, and C.
8 4.7 Stoichiometry of Precipitation Reactions 4.7 Notes Stoichiometry for reactions in solution: 1. Identify the species present in the combined solution, and determine what reaction occurs 2. Write the balanced net ionic equation for the reaction 3. Calculate the moles of reactant 4. Determine which reactant is limiting 5. Calculate the moles of product or products, as required 6. Convert to grams or other units, as required Example 4.7a Calculate the mass of solid NaCl that must be added to 1.50 L of 0.100 M AgNO 3 solution to precipitate all the Ag + ion in the form of AgCl. Example 4.7b When aqueous solutions of sodium sulfate and lead (II) nitrate are mixed, a solid is formed. Calculate the mass of the solid formed when 1.25 L of 0.0500 M lead nitrate and 2.00L of 0.0250 M sodium sulfate are mixed.
9 Worksheet 4b.02 Solution Stoichiometry Name 1. What mass of NaCl is required to make AgCl from 50.0 ml of a 0.0500 M solution of AgNO 3? 2. What volume of solid aluminum hydroxide is produced when 50.0 ml of 0.20 M aluminum nitrate is added to 200.0 ml of 0.100 M potassium hydroxide? 3. A 100.0 ml of 0.200 M potassium hydroxide is mixed with 100.0 ml of 0.200 M magnesium nitrate. Calculate the mass of solid made and the concentrations of the remaining ions in solution. 4. How many grams of silver chloride can be prepared by the reaction of 100.0 ml of 0.20 M silver nitrate with 100.0 ml of 0.15 M calcium chloride? Calculate the concentration of each remaining ion in solution after the precipitation is complete.
10 4.8 Acid-Base Reactions Read and outline 4.8 Define Acid Base Neutralization Reaction Volumetric Analysis Titration Equivalence point Stoichiometric point Indicator Endpoint What three requirements must be met for a titration to be successful? What is the most common indicator in acid-base reactions? What does standardizing the solution mean?
11 4.8 Notes We have learned that acids have H + ions and bases have OH - ions. What about ammonia, NH 3? We know that ammonia is a base but does not have hydroxide ions. In the 20s, Bronsted and Lowry came up with a definition to encompass all acids and bases. The defined an acid as a proton donor and a base as a proton acceptor. This means that NH 3 will accept a hydrogen ion from water. When was gives hydrogen (proton) all that is left is hydroxide. NH 3 + H 2 O NH 4 + + OH - Performing calculations for acid-base reaction 1. Write a balanced reaction 2. Calculate the moles of reactants. 3. Determine the limiting reactant 4. Calculate the moles of required reactant or product 5. Convert to grams or volume, as required Example 4.8a What volume of a 0.100 M HCl solution is needed to neutralize 25.0 ml of 0.350 M NaOH? Example4.8b 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. Acid-Base Titration Volumetric Analysis is a technique for determining the amount of a certain substance by doing a titration. A titration involves delivery from a buret of a measured volume of a solution of known concentration (the titrant) into a solution containing the substance being analyzed (the Analyte). The point in the titration where enough titrant has been added to react exactly with the analyte is called the equivalence point or stoichiometric point. The point where the indicator changes color is called the endpoint.
12 Three requirements for a titration to be successful 1. The exact reaction between titrant and analyte must be known. 2. The stoichiometric point must be marked accurately. 3. The volume of titrant required to reach the stoichiometric point must be known accurately. Example 4.8c A student carries out an experiment to standardize a sodium hydroxide solution. To do this, the student weighs out a 1.3009 g sample of KHC 8 H 4 O 4 known as KHP (MM = 204.22 g/mol) which 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 endpoint of the indicator. The difference between the final and initial buret reading indicates that 41.20 ml of the sodium hydroxide solution is required to react with the 1.3009 g KHP. Calculate the concentration of the sodium hydroxide solution. Example4.8d In an industrial process a chemist analyzed the waste material carbon tetrachloride and HC 7 H 5 O 2, a weak acid that has one acidic hydrogen. A sample of this waste material weighing 0.3518 g was shaken with water, and the resulting solution aqueous solution required 10.59 ml of 0.1546 M NaOH for neutralization. Calculate the mass percent of weak acid in the original sample.
13 Worksheet 4b.03 Name Acid/Base Reactions 1. What volume of each of the following acids will react completely with 50.00mL of 0.200 M NaOH? a. 0.100 M HCl b. 0.150 M HNO 3 2. What volume of each of the following bases will react completely with 25.00 ml of 0.200 M HCl? a. 0.100 M NaOH b. 0.0500 M Ba(OH) 2 3. A 25.00 ml sample of HCl solution requires 24.16 ml of 0.106 M sodium hydroxide for complete neutralization. What is the concentration of the original hydrochloric acid? 4. A student titrates an unknown amount of KHP with 20.46 ml of a 0.1000 M NaOH solution. How many grams of KHP were titrated by the sodium hydroxide solution? 5. HCl (75.0 ml of 0.250 M) is added to 225.0 ml of 0.0550 M barium hydroxide solution. What is the concentration of the excess hydrogen or hydroxide left in the solution
14 4.9 Oxidation Reduction Reactions Read and outline 4.9 Define Oxidation reduction reaction Oxidation States Oxidation Reduction Oxidizing agent Reducing agent Why is oxidation important to us? How is oxidation states assigned? What has a greater attraction for electrons than hydrogen? 4.9 Notes Oxidation Reduction Reactions are known as redox reactions. These reactions transfer electrons. In order to keep track of how many electrons are transferred, we need to assign oxidation states. Most of the time oxidation states are integer numbers. Sometimes you will have a noninteger number. Do not let this intimidate you. It is just a number. Rules for assigning oxidation numbers 2. The oxidation number for elements is zero. 3. The oxidation number for monoatomic ions is the same as its charge. 4. Fluorine in compounds is always -1. 5. Oxygen is usually assigned -2 except when in O 2-2 where it is -1, and in OF 2 where it is +2. 6. In its covalent compounds hydrogen is assigned +1. 7. The sum of the oxidation numbers should add up to the charge of the compound or polyatomic ion.
15 Examples 4.9a 1. CO 2 2. SF 6 3. NO 3-4. KMnO 4 5. NiO 2 6. K 4 Fe(CN) 6 7. (NH 4 ) 2 HPO 2 8. P 4 O 6 When balancing redox reactions not only do we need to balance the elements, but we need to balance the electrons too. After we have assigned oxidation states we will need to split the reaction into half reactions. One reaction is called the oxidation reaction. This reaction has a loss of electrons. The reactant that is being oxidized is the reducing agent. The other reaction that gains electrons is called the reduction reaction. The reactant that accepts the electrons is called the oxidizing agent. Just like we use HONCLBRIF to remember the diatomic elements, we use many ways to remember which half reaction is oxidation/reduction. There are many ways to remember this I like to say; LEO says GER. Losing electrons is oxidation and gaining electrons is reduction. Mr. Galinski uses: OIL RIG. Oxidation is losing and reduction is gaining. I like my way. You pick whatever way you are going to remember it. (My way is better!) Example4.9 b Explain the following reactions in Redox terms 2Al + 3I 2 AlI 3 2PbS + 3O 2 2PbO + 2SO 2
16 Worksheet 4b.04 Oxidation Numbers Assign the oxidation numbers to the following compounds. Name 1. Fe 3 O 4 2. XeOF 4 3. SF 4 4. CO 5. Na2C 2 O 4 6. UO 2 +2 7. As 2 O 3 8. NaBiO 3 9. As 4 10. HAsO 2 11. Mg 2 P 2 O 7 12. Na 2 S 2 O 3 13. Hg 2 Cl 2 14. Ca(NO 3 ) 2 Specify which of the following are redox reactions, and identify the oxidizing agent, the reducing agent, the substance being oxidized and the substance being reduced. 1. CH 4 + O 2 H 2 O + CO 2 2. Zn + HCl ZnCl 2 + H 2 3. Cr 2 O 7-2 + 2OH - 2CrO 4-2 + H 2 O 4. 2H 2 O 2 2 H 2 O + O 2 5. 2CuCl CuCl 2 + Cl 2
17 4.10 Balancing Oxidation and Reduction Reactions Read and outline 4.10 Define Half-Reactions What is used to balance hydrogen? What is used to balance oxygen? Why do we need to balance the electrons? 4.10 Notes Rules for balancing redox reaction in an acidic solution 1. Write oxidations states 2. Split into oxidation and reduction reactions (half-reactions) 3. Balance each reaction separately a. Balance all elements but O and H b. Balance O by adding water to the needed side c. Balance H by adding H + to the needed side d. Balance electrons 4. Add reactions together 5. Reduce. In a basic solution, add the same amount of OH - to each side of the reaction to cancel the H +.
18 Example (acid) MnO 4 - + Fe +2 Mn +2 + Fe +3 Example (acid) Cr 2 O 7-2 + C 2 H 5 OH Cr +3 + CO 2 Example (base) Ag + CN - + O 2 Ag(CN) 2 - Example (base) Al + MnO 4 - MnO 2 + Al(OH) 4 -
19 Worksheet 4b.05 Balancing Redox Reactions Name Balance the following in an acidic solution 1. Zn + H + Zn +2 + H 2 2. I - + ClO - I 3 - + Cl - 3. As 2 O 3 + NO 3 - H 3 AsO 4 + NO 4. CH 3 OH + Cr 2 O 7-2 CH 2 O + Cr +3 Balance the following in a basic solution 1. Al + MnO 4 - MnO 2 + Al(OH) 4-2. Cl 2 Cl - + OCl - 3. NO 2 - + Al NH 3 + AlO 2-4. Cr + CrO 4-2 Cr(OH) 3