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Solutions

Learn about solutions, solvents, solutes, and solubility. Discover factors that affect solubility, such as temperature and pressure. Understand solution concentration and different ways to describe it.

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Solutions

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  1. Solutions

  2. A. Definitions • Solution - homogeneous mixture containing two or more substances Solute - substance being dissolved Solvent - present in greater amount

  3. A. Definitions Solute - KMnO4 Solvent - H2O

  4. What are solutions? • Not possible to distinguish solute from solvent • Solutions can exist as gas, liquid, or solid, depending on the state of the SOLVENT • Ex: Air (gas) – oxygen in nitrogen Braces (solid) – titanium in nickel • Most solutions are liquids Homogeneous mixture

  5. Key Terms • Soluble – describes a substance that can be dissolved in a given solvent • Ex: sugar soluble in water • Insoluble – describes a substance that does not dissolve in a given solvent • Ex: sand insoluble in water

  6. Key Terms • Immiscible – describes TWO LIQUIDS that can be mixed together but separate shortly after you cease mixing them • Ex: oil and vinegar in salad dressing • Miscible – describes two liquids that are soluble in each other • Aqueous solution – solvent is water

  7. Solvation in Aqueous Solutions • Why are some substances soluble and others are not? • To form a solution, solute particles must separate from one another and the solute and solvent particles must mix.

  8. Solvation in Aqueous Solutions • Attractive forces exist between pure solute particles, between pure solvent particles, and between the solute and solvent particles. • Attractive forces between solute and solvent particles are greater than attractive forces holding solute particles together.

  9. Solvation • Solvation – the process of dissolving solute particles are surrounded by solvent particles First... solute particles are separated and pulled into solution Then...

  10. Aqueous solutions of IONIC compounds Na+ ion Cl- ion Oxygen Water molecule Hydrogen

  11. Aqueous solutions of MOLECULAR compounds • Polar molecules – molecules that have positive regions and negative regions • In molecular compounds, the negative region of the solute molecule is attracted to the positive region of water, and vice versa • Oil and water are immiscible because oil is nonpolar

  12. NONPOLAR NONPOLAR POLAR POLAR B. Solvation “Like Dissolves Like”

  13. Factors that affect rate Solubility • Agitation (stir or shake) • Temperature • Surface area

  14. Solubility • SOLUBILITY – maximum amount of solute that will dissolve in a given amount of solvent at a specified temperature and pressure • Expressed in grams of solute per 100 g of solvent

  15. Solubility Curve • shows the dependence of solubility on temperature • 170 grams of KI at 30 °C is saturated, unsaturated or supersaturated? • What temp will 100 g of sodium nitrate dissolve? • At 50 ° which has greater solubility, NaNO3 or KNO3?

  16. UNSATURATED SOLUTION more solute dissolves SATURATED SOLUTION no more solute dissolves SUPERSATURATED SOLUTION becomes unstable, crystals form Solubility concentration

  17. Supersaturated solution • To make a supersaturated solution, a saturated solution is formed at high temperature and then cooled slowly • Supersaturated solutions are unstable – if a tiny amount of solute is added, the excess solute precipitates quickly • Agitation of the container also causes crystallization

  18. Factors that affect solubility • Nature of the solute and solvent • Temperature • Affects solubility of ALL substances • Pressure • Affects the solubility of GASEOUS SOLUTES and GASEOUS SOLUTIONS

  19. Temperature and solubility • Many substances MORE soluble at HIGHER temperatures • Some substance are LESS soluble at HIGHER temperature • Gaseous solutes in liquid solvents • At higher temperature, gas particles have higher kinetic energy and escape from solution more readily

  20. Solubility of Gases

  21. Pressure and solubility • Solubility of a GAS in any solvent increases as its external pressure increases • Bottling a carbonated beverage – high pressure of carbon dioxide gas in space above liquid • Open the container, and the pressure above the liquid decreases • Bubbles of CO2 gas form in the solution, rise to top, and escape

  22. Heat of solution • Energy required to break apart solute particles and solvent particles (endothermic) • Energy released when solute and solvent particles attract each other (exothermic) • The overall energy change during the process of forming a solution is called the HEAT OF SOLUTION

  23. Solution Concentration • CONCENTRATION of a solution – a measure of how much solute is dissolved in a specific amount of solvent or solution • Qualitative descriptions • Concentrated – large amount of solute relative to its solubility • Dilute – small amount of solute relative to the solute’s solubility

  24. A. Concentration • The amount of solute in a solution. • Describing Concentration • % by mass - medicated creams • % by volume - rubbing alcohol • ppm, ppb - water contaminants • molarity - used by chemists • molality - used by chemists moles/liters Mole/Kg

  25. A. Solution Composition: Mass Percent Example: 10 grams gold in 90 grams of sand

  26. B. Solution Composition: Molarity • Concentration of a solution is the amount of solute in a given volume of solution.

  27. Quantitative Solution Concentration • Molarity (M): • Unit M read as molar • A liter of solution containing 2 moles of solute is 2M • Example: Find the molarity of a solution containing 75 g of MgCl2 in 250 mL of water. ***REMEMBER TO CONVERT mL to L

  28. Molarity • Find the molarity of a solution containing 75 g of MgCl2 in 250 mL of water. 75 g MgCl2 1 mol MgCl2 95.21 g MgCl2 0.25 L water M = moles/liters = 3.2 M MgCl2

  29. Molarity M = 75 grams / molar mass = moles 75/95.2 = .79 moles

  30. Quantitative Solution Concentration • Ex: Prepare 500 ml of a 1.54M aqueous solution of NaCl .50 L 1.54 Moles 58.5 grams 45.0 grams 1.0 L 1.0 Moles

  31. 500 mL of 1.54M NaCl 500 mL volumetric flask 500 mL mark Preparing Solutions • mass 45.0 g of NaCl • add water until total volume is 500 mL

  32. Examples • What is the molarity of .30 moles NaCl in 750 ml of water? 2. What is the molarity of 100 grams of NaCl in 2.0 liters of water? 3. How do you prepare 3 L of a 1.5M aqueous solution of HCl ? .30 moles/.750 L = 100 g/ 58.5 = 1.711.71 Moles/2.0 liters = .85

  33. Solution • Ex: Prepare 3 L of a 1.5M aqueous solution of HCl 3 L 1.5 Moles 36.5 grams 164 grams 1.0 L 1.0 Moles

  34. Dilutions • Molarity (M) • Diluting solutions • Moles of solute = molarity x liters of solution • Moles of solute before = moles of solute after dilution • Relationship before and after dilution can be written as follows: M1V1 = M2V2

  35. Dilution • What volume of 15.8M HNO3 is required to make 250 mL of a 6.0M solution? GIVEN: M1 = 15.8M V1 = ? M2 = 6.0M V2 = 250 mL WORK: M1 V1 = M2 V2 (15.8M)V1 = (6.0M)(250mL) V1 = 95 mL of 15.8M HNO3

  36. Dilution • Diluting a solution • Transfer a measured amount of original solution to a flask containing some water. • Add water to the flask to the mark (with swirling) and mix by inverting the flask.

  37. Boiling Point and Freezing Point • The presence of solute “particles” causes the liquid range to become wider. • Boiling point increases • Freezing point decreases

  38. Boiling Point and Freezing Point • Why does the boiling point of a solution increase? • Forming a bubble in a solution • Solute particles block some of the water molecules trying to enter the bubble. • Need higher pressure to maintain the bubble.

  39. Boiling Point and Freezing Point • Comparing bubbles

  40. Colligative Properties of Solutions • Colligative properties – physical properties of solutions that are affected by the number of particles but not the identity of dissolved solute particles • Vapor pressure lowering • Boiling point elevation • Freezing point depression • Osmotic pressure

  41. Ionic Substances Break into anions and cations in solution Na+ ion Cl- ion Oxygen Water molecule Hydrogen

  42. Ionic and Molecular Compounds are not the same • Molecular compounds - do not break apart when dissolving. • So, which causes a larger change in colligative properties?

  43. Chapter 16 Acids

  44. electrolytes electrolytes A. Properties ACIDS BASES • bitter taste • sour taste • turn litmus red • turn litmus blue • react with metals to form H2 gas • slippery feel • vinegar, milk, soda, apples, citrus fruits • ammonia, lye, antacid, baking soda ChemASAP

  45. Acids • Acids have Hydrogen atom in them (always put at front) • Give off – H+ • Examples: • HCl • H2SO4 • Even H2O HCl  H+ + Cl- H2SO4 2H+ + SO4-2 H2O  H+ + OH-

  46. Bases • Have OH (Called hydroxide) in them (always put at end) • Examples: • NaOH • Ca(OH)2 • Even H2O NaOH  Na+ + OH- Ca(OH)3 Ca+ + 2OH- H2O  H+ + OH-

  47. Water is amphoteric • Can act as both acid and base • Can produce both H+ • And also OH-

  48. Naming Acids • Must start with H • Two types – Binary and Oxyacids • Binary has H and one other anion • HCl • H2S • The number of H’s is equal to the charge on the anion

  49. Oxyacids • Have H and an ion from the table of common ions • Numer of H’s is equal to the charge

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