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Solution Chemistry Unit 8

Solution Chemistry Unit 8. General Chemistry Spring ’09 Mr. Hoffman. Objectives (Ch. 15). Understand and describe the basic properties of water and ice and how they effect the world around you.

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Solution Chemistry Unit 8

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  1. Solution ChemistryUnit 8 General ChemistrySpring ’09 Mr. Hoffman

  2. Objectives (Ch. 15) • Understand and describe the basic properties of water and ice and how they effect the world around you. • Explain the high surface tension and low vapor pressure of water in terms of the structure of the water molecule and hydrogen bonding (15.1.1) • Distinguish between solvent and solute (15.2.1) • Describe what happens in the solution process (15.2.2) • Explain why all ionic compounds are electrolytes (15.2.3) • Distinguish between suspension and solution (15.3.2) • Identify the distinguishing characteristic of a colloid (15.3.2)

  3. Simply Water • Ice has a low density. (Does ice float?) • It’s a polar molecule • Slightly positive (+) on one end • Slightly negative (-) on another • Look what it does to salt! • It also easily bonds to itself and easily pulls compounds apart

  4. Polar vs. Non-polar? • It’s like a tug of war… • Even pulling of electrons means it’s non-polar • Uneven pulling means it’s polar

  5. Water’s Hydrogen Bond • Water molecules are not connected by full covalent bonds, but they’re pretty strong • The formation between the hydrogen atoms on one molecule and a highly electronegative atom on another is called a hydrogen bond. • Atoms that can do this are hydrogen, oxygen, fluorine, and nitrogen • Keep reading…

  6. Hydrogen Bonds cont. • Any molecule with O-H bonds has the potential to form hydrogen bonds. • Alcohols (molecules with O-H bonds) also form hydrogen bonds. • Have similar properties to water • Proteins, nucleic acids, and carbohydrates also can form hydrogen bonds. • How they form and shape determines how they’re used biologically

  7. Water revisited • Boils at 100oC • Freezes at 0oC • Expands due to hydrogen bonding • Solid state is highly organized • One drop = ~2x1021 molecules

  8. Water in the Solid State • The structure of ice is a regular open framework of water molecules arranged like a honeycomb. • Add energy and the framework collapses • The molecules are closer together making water more dense than ice • Implications on aquatic life?

  9. Surface Tension • Surface tension: the inward force, or pull, that tends to minimize the surface area of a liquid • Causes drop to pull together • All liquids have a surface tension • Mercury has high surface tension • Surfactants • Interfere with H-bonding and reduce surface tension • Soaps and detergents

  10. Surface Tension at work

  11. Capillarity • Results from the competition between the attractive forces between the molecules of the liquid and the attractive forces between the liquid and the tube that contains it

  12. Vapor Pressure of Water • Results from the molecules escaping from the surface of the liquid and entering the vapor phase. • In water, H-bonds hold on to water molecules tightly • Tendency for molecules to escape is low • Evaporation of water is low • What would happen if it was fast?!

  13. Specific Heat of water • Specific heat: measures the amount of heat, in joules, needed to raise the temperature of 1g of substance by 1oC. • For water it’s 4.18, pretty high. • Why it takes so long to boil water • It takes a long time to absorb or release more heat for its temperature to change 1oC than a lot of other substances. • Think of a pool in the summer time.

  14. Specific Heat

  15. Water: The Universal Solvent • Almost always found in solution • A very good solvent due to its polar abilities • Examples of aqueous solutions • Milk • Soda pop • Coffee and tea • Tap water • Look at the ingredient list of a liquidy beverage. Water is probably there.

  16. Solutions • When one substance dissolves into another, that is called a SOLUTION • Example: sugar water, Kool-Aid • There are two main parts of a solution: • SOLUTE= the dissolved material • Example: sugar, salt, oxygen (air) • SOLVENT= the substance that is doing the dissolving (usually a liquid) • Usually present in the highest amount • Example: Water, nitrogen (air)

  17. Dissolving • Water, a polar molecule, is capable of dissolving a range of compounds • Many ionic compounds (like NaCl) are soluble in water • When dissolved, ionic solutions are very good conductors of electricity

  18. Electrolytes vs. Nonelectrolytes • Electrolyte • A compound that conducts an electric current when it is in an aqueous solution • All ionic compounds are electrolytes because they dissociate into ions • Nonelectrolyte • A compound that does not conduct an electric current in aqueous solution • Many covalents are this because they are not composed of ions

  19. How does dissolving happen? • Ionic solids are composed of positive and negative ions. • Water has a positive and negative end (it’s polar) • Opposites attract and the ionic compounds separate into ions. • The process by which charged particles in an ionic solid separate from one another is solvation

  20. Dissolving Covalent substances • Sugar is the best example • Almost 200 grams can dissolve in 100 mL of H2O! • It has O-H bonds, which makes it polar, so it’s easily dissolvable in water • If a molecule contains O-H bonds, it will tend to be polar and it can form hydrogen bonds with water.

  21. Dissolving Covalent substances • Covalent molecules are simply separated from one another by water molecules. • They don’t solvate into separate ions • Covalent solutions can’t conduct electricity

  22. “Like Dissolves Like” • This means that dissolving occurs when similarities exist between the solvent and the solute. • Sugar is a polar molecule, so is water, and water tends to dissolve substances that are polar or that form hydrogen bonds. • Oil and water don’t mix. • Oil is nonpolar • But different oils are “like” enough to mix and stay mixed.

  23. Suspensions • A (heterogeneous) mixture from which particles settle out upon standing • A suspension differs from a solution because the particles are much larger and do not stay suspended indefinitely • Cornstarch mixed with water thickens sauces

  24. Suspensions • Two substances are clearly identified • Dispersed phase • Ex) clay, dirt, sand, flour • Dispersion medium • Water, ethanol • Think of a glass of water with sand or mud in it. • Typically easy to separate

  25. Colloids • A heterogeneous mixture containing particles that range in size from 1nm to 1000nm • Particles spread throughout the dispersion medium (s, l, g) • Glues • Gelatin • Paint • Aerosol sprays • smoke

  26. Colloid Examples (Table 15.3)

  27. Tyndall Effect, Coagulation • Tyndall Effect • The scattering of visible light by colloidal particles • Suspensions also do this but solutions don’t. • Coagulation • The clumping of particles in a colloid

  28. Emulsions • A colloidal dispersion of a liquid in a liquid • Must have an “emulsifying agent” • To form the emulsion • To maintain stability • Ex) soap, detergent, egg yolk

  29. Objectives (Ch. 16) • Identify the factors that determine the rate at which a solute dissolves (16.1.1) • Identify the units usually used to express the solubility of a solute (16.1.2) • Identify the factors that determine the mass of solute that will dissolve in a given mass of solute (16.1.3)

  30. Solution Formation • Determining factors • Composition of solvent • Composition of the solute • Speed of dissolving factors • Stirring (agitation) • Temperature • Surface area • All involve contact between solvent and solute

  31. Rate of Dissolving • Stirring the Solution • Increases the interaction between water molecules and the solute. • Solute and solvent interact more often, the rate of dissolution is faster. • Does not influence the amount of solute that will dissolve • Oil will never mix with water not matter how long you stir or shake that Italian dressing

  32. Rate of Dissolving • Heating the Solution • Increases kinetic energy of the water molecules • Increases frequency and force of the collisions between solute and solvent

  33. Rate of Dissolving • Grinding the solute • Creates more surface area (remember the big fireball demo?) • Solvent molecules attack the edged surfaces of solute crystals. • The more surface area expose, the faster the rate of dissolving

  34. Solubility • When solute enters the solvent… • Particles move from the solid into the solution • Other dissolved particles move from solution back to the solid • Occurs at the same rate • Called a saturated solution • Will stay this way as long as temperature stays the same

  35. Solubility • The solubility of a substance is the amount of solute that dissolves in a given amount of solvent at a specified temperature and pressure to produce a saturated solution • Units: grams per liter (g/L) • Miscible • two liquids that dissolve in each other • Immiscible • Two liquids not soluble in each other

  36. Types of Solutions • Saturated solution • Solution holding the max. amt. Of solute per amt. Of solution under given conditions. • Add more solute it won’t dissolve • Unsaturated solution • The amt. of solute is less than the max that could be dissolved. • Add more solute it will dissolve

  37. Solutions • Supersaturated solution • Contain more solute than the usual max. amt. And are unstable. • Add a crystal and it fills the container with crystals

  38. Effects of Pressure • Huge effect on gases, very little on solids and liquids • Gas solubility increases as the partial pressure of the gas above the solution increases. (direct relationship) • Ex) Soda bottle has lots of dissolved CO2 in it which is forced in at the plant. • When you open the bottle you hear a hiss and CO2 starts escaping from the bottle decreasing the concentration on CO2 in the bottle

  39. Solubility Curves • The solubility of substances changes with temperature • For example, is it easier to dissolve sugar in hot or cold coffee? • Solids become more soluble at higher temperatures • Gases become less soluble at higher temperatures

  40. Solubility Curves (cont.) • Scientist have studied many substances solubility at different temperatures • They created graphs which show this data

  41. Solubility Curves (cont.) • Let’s simplify the graph with all the substances down to just one substance

  42. Solubility Curves (cont.) • What does this graph tell you about KCl at 80°C? • 52g of KCl dissolve in 100g of water Is KCl a solid or gas in this graph?

  43. Solubility Curves (cont.) • How many grams of KCl will dissolve in 500g of water at 80°C? • 260g of KCl (52g x 5 = 260g)

  44. Solubility Curves (cont.) • How many grams of water will it take to dissolve 26 g of KCl at 80°C? • 50g of H2O (1/2 of what dissolves in 100g H2O) (% of 100g: 26g/52g=.50)

  45. Solubility Curves (cont.) • If one dissolves 95 grams of KCl in 250 grams of water at 80°C, what kind of solution will they have? • Unsaturated You need to determine the saturation point before you can decide the type of solution. (Sat. pt. From graph)x(%H2O) 52 g x 2.5 = 130 g Saturation point in 250g of water

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