Substances, Compounds & Mixtures

1. Substances, Compounds & Mixtures How everything is put together.

2. Substances • Matter that has the same composition and properties throughout is called a substance. • When different elements combine, other substances are formed.

3. Substances • Contains only one particle • Can exist in 3 states of matter • Can be elements or compounds Picture from http://www.ilpi.com/msds/ref/gifs/statesofmatter.gif

4. Mixtures • A mixture is a combination of two or more substances where there is no chemical combination or reaction.

5. A mixture is a combination of two or more substances where there is no chemical combination or reaction.

6. Mixtures combine physically in no specific proportions. They just mix.

7. Solids, liquids and gases can be combined to create a mixture.

8. Mixture Types • MIXTURES MAY BE HETEROGENEOUS OR HOMOGENEOUS

9. Heterogeneous Mixtures: • The prefix: "hetero"- indicates difference • A heterogeneous mixture consists of visibly different substances or phases • Two or more parts can be seen

10. Examples: • Pizza • Sandwich • Chex Mix

11. Suspensions • A SUSPENSION is a heterogeneous mixture of large particles • These particles are visible and will settle out on standing • Examples of suspensions are: fine sand or silt in water or Italian salad dressing

12. Homogeneous Mixtures • Homogeneous Mixtures: • The prefix: "homo"- indicates the same • Have the same uniform appearance and composition throughout

13. Solutions • SOLUTIONS are homogeneous mixtures

14. What is a solution? • A solution is a mixture of two or more substances. • At least two substances must be mixed in order to have a solution

15. A solution has two parts • The substance in the larger amount is called the SOLVENT - it does the dissolving • IN most common instances water is the solvent • The substance in the smallest amount and the one that DISSOLVES is called the SOLUTE

16. Examples of solutions • Salt water • Clean Air • Vinegar

17. Types of solutions Solute Solvent Example Metals dissolved in metals are called alloys.

18. Solution Salt water is considered a solution. How can it be physically separated? • a mixture of two or more substances that is identical throughout • can be physically separated • composed of solutes and solvents the substance in the smallest amount and the one that dissolves in the solvent the substance in the larger amount that dissolves the solute Iced Tea Mix (solute) Water (solvent) Iced Tea (solution) Colloids (milk, fog, jello) are considered solutions

19. Solutes Change Solvents • The amount of solute in a solution determines how much the physical properties of the solvent are changed • Examples: Lowering the Freezing Point Raising the Boiling Point The freezing point of a liquid solvent decreases when a solute is dissolved in it. Ex. Pure water freezes at 320F (00C), but when salt is dissolved in it, the freezing point is lowered. This is why people use salt to melt ice. The boiling point of a solution is higher than the boiling point of the solvent. Therefore, a solution can remain a liquid at a higher temperature than its pure solvent. Ex. The boiling point of pure water is 2120F (1000C), but when salt is dissolved in it, the boiling point is higher. This is why it takes salt water longer to boil than fresh water.

20. Concentration • the amount of solute dissolved in a solvent at a given temperature • described as dilute if it has • a low concentration of • solute • described as saturated if it • has a high concentration of • solute • described as supersaturatedif • contains more dissolved solute • than normally possible

21. Solubility • the amount of solute that dissolves in a certain amount of a solvent at a given temperature and pressure to produce a saturated solution • influenced by: What do we call things that are not soluble? Pressure Temperature Solids increased temperature causes them to be more soluble and vice versa Gases increased temperature causes them to be less soluble and vice versa Ex. Iced Coffee Solids increased pressure has no effect on solubility Gases increased pressure causes them to be more soluble and vice versa Ex. Soda, “The Bends”

22. Solutions How does a solid dissolve into a liquid? What ‘drives’ the dissolution process? What are the energetics of dissolution?

23. How Does a Solution Form? • Solvent molecules attracted to surface ions. • Each ion is surrounded by solvent molecules. • Enthalpy (DH) changes with each interaction broken or formed. Ionic solid dissolving in water

24. How Does a Solution Form? • Solvent molecules attracted to surface ions. • Each ion is surrounded by solvent molecules. • Enthalpy (DH) changes with each interaction broken or formed.

25. How Does a Solution Form The ions are solvated (surrounded by solvent). If the solvent is water, the ions are hydrated. The intermolecular force here is ion-dipole.

26. dry Dissolution vs reaction NiCl2(s) Ni(s) + HCl(aq) NiCl2(aq) + H2(g) • Dissolution is a physical change—you can get back the original solute by evaporating the solvent. • If you can’t, the substance didn’t dissolve, it reacted.

27. Degree of saturation • Saturated solution • Solvent holds as much solute as is possible at that temperature. • Undissolved solid remains in flask. • Dissolved solute is in dynamic equilibrium with solid solute particles.

28. Degree of saturation • Unsaturated Solution • Less than the maximum amount of solute for that temperature is dissolved in the solvent. • No solid remains in flask.

29. Degree of saturation • Supersaturated • Solvent holds more solute than is normally possible at that temperature. • These solutions are unstable; crystallization can often be stimulated by adding a “seed crystal” or scratching the side of the flask.

30. Degree of saturation Unsaturated, Saturated or Supersaturated?  How much solute can be dissolved in a solution? More on this in Chap 17 (solubility products, p 739)

31. Factors Affecting Solubility • Chemists use the axiom “like dissolves like”: • Polar substances tend to dissolve in polar solvents. • Nonpolar substances tend to dissolve in nonpolar solvents.

32. Factors Affecting Solubility The stronger the intermolecular attractions between solute and solvent, the more likely the solute will dissolve. Example: ethanol in water Ethanol = CH3CH2OH Intermolecular forces = H-bonds; dipole-dipole; dispersion Ions in water also have ion-dipole forces.

33. Factors Affecting Solubility Glucose (which has hydrogen bonding) is very soluble in water. Cyclohexane (which only has dispersion forces) is not water-soluble.

34. Temperature Generally, the solubility of solid solutes in liquid solvents increases with increasing temperature.

35. Temperature • The opposite is true of gases. Higher temperature drives gases out of solution. • Carbonated soft drinks are more “bubbly” if stored in the refrigerator. • Warm lakes have less O2 dissolved in them than cool lakes.

36. mol of solute L of solution M = Molarity (M) • You will recall this concentration measure from Chapter 4. • Because volume is temperature dependent, molarity can change with temperature.

37. mol of solute kg of solvent m = Molality (m) Because neither moles nor mass change with temperature, molality (unlike molarity) is not temperature dependent.

38. moles of A total moles in solution XA = Mole Fraction (X) • In some applications, one needs the mole fraction of solvent, not solute—make sure you find the quantity you need!

39. Colligative Properties

40. Colligative Properties • All colligative properties • Depend on the number and not the nature of the solute molecules • Due to reduction in chemical potential in solution vs. that of the pure solvent • Freezing point depression • Boiling Point Elevation • Osmotic Pressure

41. Boiling Point Elevation Examine the chemical potential expressions involved

42. Boiling Point Elevation #2 The boiling point elevation

43. Freezing Point Depression Examine the chemical potential expressions involved

44. Freezing Point Depression #2 Define the freezing point depression

45. Osmosis

46. Osmosis • The movement of water through a semi-permeable membrane from dilute side to concentrated side • the movement is such that the two sides might end up with the same concentration • Osmotic pressure: the pressure required to prevent this movement

47. Osmosis – The Thermodynamic Formulation  - the osmotic pressure Equilibrium is established across membrane under isothermal conditions

48. The Final Equation The osmotic pressure is related to the solutions molarity as follows

49. Terminology Isotonic: having the same osmotic pressure Hypertonic: having a higher osmotic pressure Hypotonic: having a lower osmotic pressure

50. Terminology #2 Hemolysis: the process that ruptures a cell placed in a solution that is hypotonic to the cell’s fluid Crenation: the opposite effect