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Diluting Solutions

Diluting Solutions. You can make a less concentrated solution of a known solution by adding a measured amount of additional solvent to the standard solution. The number of molecules, or moles of solute remains the same before and after the dilution. Side note on formulas for this module.

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Diluting Solutions

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  1. Diluting Solutions • You can make a less concentrated solution of a known solution by adding a measured amount of additional solvent to the standard solution. • The number of molecules, or moles of solute remains the same before and after the dilution

  2. Side note on formulas for this module • There are no formulas given on the test!! • You must MEMORIZE the formulas, so start doing so NOW so it won’t be an issue come test time! • The test is similar to the last one in that there are lots of calculations, you get MARKS for showing your work, knowing formulas, etc. • This is definitely the heaviest module yet, so stay on top of it.

  3. New Formula!! • For dilution problems use the formula: • *V2 is the total final volume of the concentrated solution plus the water added

  4. Sample Problem • What is the concentration of the solution that is obtained by diluting 300mL of 6.00mol/L HCI solution to a final volume of 900mL? • Given? • Required?

  5. Sample Problem • Suppose you are given a solution of 1.25mol/L NaCl(aq). What volume must you dilute to prepare a 50mL of 1.00mol/L NaCl(aq) solution?

  6. Sample Problem • Perform the calculation and outline the laboratory procedure required to prepare 500mL of 0.80mol/L KNO3 solution from a concentrated 4.00mol/L KNO3 solution.

  7. Procedure figure 8.22 page 319 • Measure 100mL of the concentrated 4.00mol/L KNO3 solution into a graduated cylinder. • Carefully pour, using a funnel, into a 500mL volumetric flask • Rinse the graduated cylinder and add the rinsing to the flask • Add water to the flask until almost full • Add the water drop by drop until the bottom of the meniscus touches the line on the neck of the flask. • Label the flask 0.8mol/L KNO3 and date it.

  8. Aqueous Solutions • In general: • “Soluble” means that more than: • “Insoluble” means that the solubility:

  9. How can you tell if something is soluble, slightly soluble or insoluble? (Review)

  10. Factors that affect solubility of ionic substances: • The effect of ion charge • The effect of ion size

  11. 1. Ion Charge • Compounds with • Compounds with • Why? • Ex: phosphates (PO4-3) compounds tend to be insoluble and alkali metals compounds (Na+, K+) tend to be soluble

  12. 2. Ion size • Gaining or losing e- changes the size of the atoms Why? Think about energy levels! Bohr-Rutherford Diagrams!

  13. Cations

  14. Anions

  15. Bonding • Bonds between small ions are closer then large ions with the same charge • Thus, • This in turn affect solubility:

  16. Example-Halogens Which is more soluble? NaF NaCl NaBr NaI LiF NaF KF CsF

  17. Sulfides and Oxides • Influenced by size and charge • Tend to be insoluble because of relatively small size and -2 charge

  18. Using a Solubility Table

  19. General Solubility Guidelines • Higher guidelines always take precedence • Eg. BaCl2 • Ba is listed as GL 4 and insoluble • Cl is listed as GL 3 and soluble • The higher GL number takes precedence so you can assume that

  20. Soluble or Insoluble? • A) Lead (II) chloride, white crystalline powder used in paints • B) Zinc oxide, white pigment used in paints, cosmetics, and calamine lotion • C) Silver acetate, AgCH3COO, white powder used to help people quit smoking because of it’s bitter taste

  21. Reactions in Aqueous Solutions • When mixing two aqueous ionic compounds together, there are 2 possible outcomes: • 1. • 2. • We have seen this before • Remember? Lead (II) nitrate and potassium iodide

  22. Double Displacement Reactions • Formation of a gas • Formation of a precipitate • Formation of water (H and OH ions are removed from solution as water)

  23. Precipitation Reactions • A precipitate is an insoluble substance • How can we predict if there will be a precipitation reaction without doing it in a lab? • Think about the exchange of ions first • Then use the solubility guidelines • Eg. Lead (II) nitrate and potassium iodide

  24. Precipitation Reactions • Pb(NO3)2(aq) + 2KI(aq)  2KNO3(aq) + PbI2 • Looking at the products and the solubility guidelines • Potassium salts and nitrates are soluble • Pb2+ is insoluble, GL 2 • I- is soluble, GL 3

  25. Sample Problems • Potassium carbonate + copper (II) sulfate • Ammonium chloride + zinc sulfate • Hint: you need to get the formulas, do a double displacement, use the solubility guidelines to find if there will be a precipitate

  26. I need a BREAK! YOUR turn!! • Break into three groups • Group 1 – pg. 357 • Group 2 – pg. 358 • Group 3 – pg. 360 (top only) • Summarize the important info on the page and present it to the class • Make sure to write down points on the board so we can copy them down!

  27. Homework • Self Check #2 • Readings for Lesson 7 – Acids and Bases • Try the questions in the green book for Lesson #7 • Print off Lab sheet and complete Pre-Lab questions • Make sure to come to the LAB on Monday • Re-writes 4:30 tomorrow

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