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Parts of Solutions

Parts of Solutions. Solution- homogeneous mixture. Solute- what gets dissolved. Solvent- what does the dissolving. Soluble- Can be dissolved. Miscible- liquids dissolve in each other. Figure 4.1 The Water Molecule. Hydration. The process of breaking the ions of salts apart.

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Parts of Solutions

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  1. Parts of Solutions • Solution- homogeneous mixture. • Solute- what gets dissolved. • Solvent- what does the dissolving. • Soluble- Can be dissolved. • Miscible- liquids dissolve in each other.

  2. Figure 4.1 The Water Molecule

  3. Hydration • The process of breaking the ions of salts apart. • Ions have charges and are attracted to the opposite charges on the water molecules.

  4. Figure 4.2 Polar Water Molecules Interact with the Positive and Negative Ions of a Salt Assisting in the Dissolving Process

  5. H H H H O O O H H H H O O H H O O H H H H H H O H O H How Ionic solids dissolve Click here for Animation

  6. Solubility • How much of a substance will dissolve in a given amount of water. • Usually g/100 mL • Varies greatly, but if they do dissolve the ions are separated, • and they can move around. • Water can also dissolve non-ionic compounds if they have polar bonds.

  7. Figure 4.3a The Ethanol Molecule Contains a Polar O-H Bond Similar to Those in the Water Molecule

  8. Figure 4.3b The Polar Water Molecule Interacts Strongly with the Polar-O-H bond in Ethanol

  9. Electrolytes • Electricity is moving charges. • The ions that are dissolved can move. • Solutions of ionic compounds can conduct electricity. • Solutions are classified three ways.

  10. Types of solutions • Strong electrolytes- completely dissociate (fall apart into ions). • Many ions- Conduct well. • Weak electrolytes- Partially fall apart into ions. • Few ions -Conduct electricity slightly. • Non-electrolytes- Don’t fall apart. • No ions- Don’t conduct.

  11. Figure 4.4a-c Electrical Conductivity of Aqueous Solutions

  12. Figure 4.5 NaCl Dissolves

  13. Figure 4.6 HCL is Completely Ionized

  14. Figure 4.7 An Aqueous Solution of Sodium Hydroxide

  15. Figure 4.8 Acetic Acid (HC2H3O2)

  16. Figure 4.9 The Reaction of NH3 in Water

  17. Measuring Solutions • Concentration- how much is dissolved. • Molarity = Moles of solute Liters of solution • abbreviated M • 1 M = 1 mol solute / 1 liter solution • Calculate the molarity of a solution with 34.6 g of NaCl dissolved in 125 mL of solution.

  18. Figure 4.10a-c Steps Involved in the Preparation of a Standard Aqueous Solution

  19. Figure 4.11a-b Measuring Pipets and Volumetric Pipets Measure Liquid Volume

  20. Figure 4.12a-c A Measuring Pipet is Used to Add Acetic Solution to a Volumetric Flask

  21. Figure 4.14 a&b Reactant Solutions

  22. Figure 4.15 a&b The Reaction of K2CrO4 and Ba(NO3)2

  23. Figure 4.17 Molecular-Level Representations Illustrating the Reaction of KCl (aq) with AgNO3 (aq) to Form AgCl (s)

  24. Writing Net Ionic Equations • Strong acids are all written in net ionic form • -Binary acids – are all strong (except for HF(aq) ) • -Oxyacids-If the number of oxygens exceeds the number of hydrogens by 2 or more they are considered strong. • -Polyprotic acids ionize one (1) hydrogen at a time. All subsequent ionizations of acidic hydrogens are considered weak. • (except for HSO4(aq)1- ) • H2SO4(aq) → H(aq)1+ + HSO4(aq)1- • Strong bases are all written in ionic form. • -Group IA and IIA metal hydroxides are strong bases. • All weak acids and bases (those not mentioned above) are always written in molecular form. • Ionic Salts-If soluble-written in ionic form • -if insoluble written in molecular/undissociated form.

  25. Writing Net Ionic Equations • Oxides are always written in molecular/undissociated form. • Gases are always written in molecular form. • Molecular compounds are always written in molecular form.

  26. Determining the Mass of Product Formed

  27. Determining the Mass of Product Formed

  28. Performing Calculations for Acid-Base Reactions

  29. Neutralization Reactions I

  30. Neutralization Reactions II

  31. Neutralization Titration

  32. Figure 4.19 The Reaction of Solid Sodium and Gaseous Chlorine to Form Solid Sodium Chloride

  33. Figure 4.20 A Summary of Oxidation-Reduction Process

  34. The Half-Reaction Method (Acidic Solution)

  35. The Half-Reaction Method (Basic Solution)

  36. Figure 4.4a-c Electrical Conductivity of Aqueous Solutions

  37. An Aqueous Solution of Co(NO3)2.

  38. Figure 4.10 Steps Involved in the Preparation of a Standard Aqueous Solution

  39. Figure 4.13 Yellow Aqueous Potassium

  40. Figure 4.14a-b Reactant Solutions

  41. Figure 4.15c Solution Post-Reaction

  42. Figure 4.16 Addition of Silver Nitrate to Aqueous Solution of Potassium Chloride

  43. Figure 4.17 Reaction of KCI(aq) with AgNO3(aq) to form AgCI(s).

  44. Lead Sulfate

  45. KOH and Fe(NO3)3 Mix to Create Solid Fe(OH)3.

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