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AQUEOUS SOLUTIONS. Solute Concentration Molecular Weight = Sum of weight of all atoms in a molecule (expressed in Daltons). For example: Determine a mole of CH 3 COOH CH 3 COOH 2 C 2 x 12 Da = 24 4 H 4 x 1 Da = 4

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  1. AQUEOUS SOLUTIONS Solute Concentration Molecular Weight = Sum of weight of all atoms in a molecule (expressed in Daltons). For example: Determine a mole of CH3COOH CH3COOH 2 C 2 x 12 Da = 24 4 H 4 x 1 Da = 4 2 O 2 x 16 Da = 32 --------------------- M.W. Da = 60 g/mol

  2. AQUEOUS SOLUTIONS Mole = Amount of a substance that has a mass in grams numerically equivalent to its molecular weight in Daltons. For example: To determine a mole of sucrose (C12H22O11). Calculate molecular weight: C = 12 Da 12 Da x 12 = 144 Da H = 1 Da 1 Da x 22 = 22 Da O = 16 Da 16 Da x 11 = 176 Da 342 Da Express it in grams (342 g).

  3. AQUEOUS SOLUTIONS Molarity = Number of moles of solute per liter of solution. For example: To obtain 1 M sucrose solution, weigh out 342 g and add water up to 1L. Advantage of measuring in moles: 1. Rescales weighing single molecules in Daltons to grams, which is more practical for laboratory use. 2. A mole of one substance has the same number of molecules as a mole of any other substance (6.02 x 1023; Avogadro’s number). 3. Allows one to combine substances in fixed ratios of molecules

  4. AQUEOUS SOLUTIONS For example: How much CH3COOHdo you need to make 1 liter of a 2 M solution of CH3COOH? Ask yourself these questions: • How many grams of this substance are in 1 mole of CH3COOH? 2. How many liters are you trying to prepare? 3. What is the ultimate concentration you are trying to prepare (M)?

  5. AQUEOUS SOLUTIONS Answer: 60 grams/mole x 1 liter x 2 moles/liter = 120 g of CH3COOH is needed to prepare a 2 M solution.

  6. AQUEOUS SOLUTIONS For Example: How much CaCl2do you need to make 750 ml of a 1.5 M CaCl2solution? How many grams/mole? 110 g/mol How many liters? .750 L What is M? 1.5 M

  7. AQUEOUS SOLUTIONS Answer: 110 g/mole x .750 L x 1.5 moles/liter = 123 g You will need 123g of CaCl2

  8. AQUEOUS SOLUTIONS For Example: You want to make a 10 M C6H12O6 solution. If you have 900 grams of C6H12O6, how much can you make? How many grams/mole? 180 g/mol What is the M desired? 10 M How many grams total do you have to work with? 900 grams

  9. AQUEOUS SOLUTIONS Answer: You can make 0.5 L of a 10 M C6H12O6 solution

  10. AQUEOUS SOLUTIONS What if the question would have asked for a 2.5 M C6H12O6 solution? What about a 4 M C6H12O6 solution?

  11. AQUEOUS SOLUTIONS Answer: 5 moles / 2.5 moles/liter = 2 L 5 moles / 4 moles/liter = 1.25 L

  12. AQUEOUS SOLUTIONS For Example: What is the molar concentration of a NaOH solution where you have been given 60 grams of NaOH and ask to prepare a final volume of 3 liters? What is the amount of grams/mole? = 40 g/mole How many grams do you have? 60 grams What is your desired volume? 3 L

  13. AQUEOUS SOLUTIONS Answer: 60 grams / 40 grams/mole = 1.5 mole 1.5 moles / 3 liters = 0.5 moles/liter The molar concentration is 0.5 M.

  14. Acids, Bases and pH Dissociation of Water Molecules (pH) Occasionally, the hydrogen atom that is shared in a hydrogen bond between two water molecules, shifts from the oxygen atom to which it is covalently bonded to the unshared orbitals of the oxygen atom to which it is hydrogen bonded.

  15. Figure 3.9 The pH of some aqueous solutions

  16. Acids, Bases and pH 1. Only a hydrogen ion (proton with +1 charge) is actually transferred. 2. Transferred proton binds to an unshared orbital of the second water molecule creating a hydronium ion (H3O+). 3. Water molecule that lost a proton has a net negative charge and is called a hydroxide ion (OH-). H2O + H2O ↔ H3O+ + OH-

  17. Unnumbered Figure (page 47) Chemical reaction: hydrogen bond shift

  18. Acids, Bases and pH 4. By convention, ionization of H2O is expressed as the dissociation into H+ and OH- H2O ↔ H+ + OH- 5. Reaction is reversible. 6. At equilibrium, most of the H2O is not ionized.

  19. Acids, Bases and pH 1. At equilibrium in pure water at 25oC: 2. Number of H+ ions + number of OH- ions. [H+] = [OH-] = 1____ M = 10 –7 M 10,000,000 *Note that brackets indicate molar concentration*

  20. Acids, Bases and pH A solution in which: [H+] = [OH-] is a neutral solution. [H+] > [OH-] is an acidic solution. [H+] < [OH-] is a basic solution.

  21. Acids, Bases and pH The pH Scale In any aqueous solution: [H+] [OH-] = 10-14 M2 In a neutral solution, [H+] =10-7 M and [OH-] =10-7 M. In an acidic solution if the [H+] = 10-5 M, then [OH-] =10-9 M. In a basic solution if the [H+] =10-9 M, then [OH-] = 10-5 M.

  22. Acids, Bases and pH pH = Negative log10 of the [H+] expressed in moles per liter. pH of 7 is a neutral solution. pH < 7 is an acidic solution. PH > 7 is a basic solution. Most biological fluids are within the pH range of 6 to 8. There are some exceptions such as stomach acid with pH of 1.

  23. Acids, Bases and pH Each pH unit represents a tenfold difference (scale is logarithmic ), so a slight change in pH represents a large change in actual [H+]. pH = -log [H+] or [H+] = 10 – pH M pOH = -log [OH-] or [OH-] = 10 – pOH M **pH + pOH = 14**

  24. Acids, Bases and pH For Example : If the concentration of OH- in an aqueous solution is 10-3, what is the pH?

  25. Acids, Bases and pH pOH = -log [OH-] pH + pOH = 14 pOH = -log [10-3] pH + 3 = 14 pOH = 3 pH = 14 – 3 Final Answer pH = 11

  26. Acids, Bases and pH For Example: What is the H+ concentration in a solution that has a pH of 7?

  27. Acids, Bases and pH Answer: 1 X 10 –7 M

  28. Acids, Bases and pH For example: How much greater is the [H+] in a solution with pH 2 than in a solution with pH 6?

  29. Acids, Bases and pH Answer: pH 2 = [H+] of 10-2= 1 M 100 pH 6 = [H+] of 10-6 = 1 M 1,000,000 10,000 times greater.

  30. Acids, Bases and pH Buffers By minimizing wide fluctuations in pH, buffers help organisms maintain the pH of body fluids within the narrow range necessary for life (usually pH 6-8).

  31. Acids, Bases and pH Buffer 1. Substances that prevent large sudden changes in pH. 2. Are combinations of H+ -donor and H+ -acceptor forms of weak acids or bases. 3. Work by accepting H+ ions from solution when they are in excess, and by donating H+ ions to the solution when they have been depleted.

  32. Acids, Bases and pH For example: Bicarbonate buffer H2CO3 HCO3- + H+ H+ donor H+ acceptor Weak acid Weak base

  33. Acids, Bases and pH HCl + NaHCO3 H2CO3 + NaCl Strong acid Weak acid NaOH + H2CO3 NaHCO3 + H2O Strong base Weak base

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