States of Matter

1. States of Matter • For molecular substances there are basically three states or phases of matter. • Solids (s): Molecules are held in place by intermolecular interactions. • Liquids (l): Molecules are held next to one another by noncovalent, intermolecular interactions, however, these interactions are not strong enough to prevent the molecules from flowing past one another. • Gases (s): The intermolecular interactions are too weak to hold the molecules next to one another, so the molecules wander off on their own.

2. States of Matter Water vapor, H2O(g) Ice, H2O(s) Liquid water, H2O(l)

3. States of Matter • The strength and numbers of the noncovalent intermolecular interactions determine which state a molecular substance is in. • The predominant noncovalent interaction between water molecules is the hydrogen bond:

4. States of Matter • These interactions can be disrupted by adding heat. • Adding heat increases the kinetic energy of the molecules • This is most readily observed with gases by looking at the Ideal Gas Law equation: • As the temperature of a gas increases, so does its kinetic energy.

5. States of Matter • Ideal Gas Law Simulation

6. States of Matter • A as heat is added to a molecular substance, it warms until reaching one of the phase transition temperature. • At that point the heat (kinetic energy) that is added to the substance is used to break the noncovalent, intermolecular interactions.

7. States of Matter • For a more detailed description of phase transitions, along with an animation of the process,see the Chem 150Elaboration - States of Matter

8. Enthalpy, Entropy and Free Energy • Energy is defined as the ability to do work. • The heat energy we have been talking about is also called Enthalpy (H) • It was used to do the work of breaking the noncovalent, intermolecular interactions present in solids and liquids. • When Enthalpy is put into an object, such as an ice cube, the change in Enthalpy for the ice cube increases. • (ΔH > 0). the Δ symbol means “change in”. • Changes in nature can be either spontaneous (favorable), or nonspontaneous (unfavorable).

9. Enthalpy, Entropy and Free Energy • Why are some changes spontaneous while others are nonspontaneous? • Why, like the ice on a pond, are changes spontaneous some of the time and nonspontaneous at other times? • Asking some questions about the energy changes that take place can to help answer theses questions

10. Enthalpy, Entropy and Free Energy • Changes occur spontaneously in nature when energy is released. • The case of the rolling stone.

11. Enthalpy, Entropy and Free Energy • Although Enthalpy is a form of energy, it alone cannot be used to answer these questions. • The melting of the ice from a pond on a warm spring day is spontaneous • However, the ice is absorbing heat (ΔΗ > 0) (endothermic) • A second factor called Entropy (S), needs to also be considered to determine if a change is spontaneous or nonspontaneous.

12. Enthalpy, Entropy and Free Energy • Entropy is a measure of disorder. • When ΔS > 0, things become more disorder. • Nature prefers things to be disordered:

13. Enthalpy, Entropy and Free Energy • Enthalpy and Entropy can be combined to calculate another type of energy called Free Energy (G). • ΔG = ΔΗ - ΤΔS • The change in Free Energy can be used to predict whether a change is spontaneous or nonspontaneous. • When ΔG < 0, the change is spontaneous (favorable) • When ΔG > 0, the change is nonspontaneous (unfavorable)

14. Enthalpy, Entropy and Free Energy • When Ice melts, ΔH > 0 and ΔS > 0 • It gains heat and becomes more disordered • Above the freezing temperature, TΔS > ΔH and ΔG is negative ( ΔG < 0) • Ice melts spontaneously. • Below the freezing Temperature, TΔS < ΔH and ΔG is positive ( ΔG > 0) • Ice does not melt spontaneously.

15. Enthalpy, Entropy and Free Energy • When Ice freezes, ΔH < 0 and ΔS < 0 • It loses heat and becomes more ordered • ΔH makes a negative contribution to ΔG • ΔS makes a positive contribution to ΔG • Below the freezing Temperature, the magnitude of TΔS < ΔH and ΔG is NEGATIVE (ΔG < 0) • Ice freezes spontaneously. • Above the freezing temperature, the magnitude of TΔS > ΔH and ΔG is positive (ΔG > 0) • Ice freezes nonspontaneously.

16. Enthalpy, Entropy and Free Energy • For a more detailed description using the enthalpy, entropy an free energy changes to predict if a process is spontaneous or not,see the Chem 150Elaboration - Ethalpy, Entropy & Free Energy

17. Enthalpy, Entropy and Free Energy • Figure 5.6, Raymond ΔS < 0 (molecules are more ordered) ΔS > 0 (molecules are more disordered)

18. Liquids • Liquids have various physical properties that reflect the strength of the intermolecular interactions that hold the liquid together • Boiling point temperature • Viscosity • Resistance to flow • Vapor pressure

19. Liquids • Viscosity • Resistance to flow

20. 1 atm = 760 Torr Liquids • Vapor Pressure and Boiling Points are related • The boiling point is the temperature at which the vapor pressure is equal to the atmospheric pressure.

21. 1 atm = 760 Torr Liquids • Vapor Pressure and Boiling Points are related • The boiling point is the temperature at which the vapor pressure is equal to the atmospheric pressure.

22. Questions (Clickers) • You planning to do some surgery on your kitchen table and know that you need to sterilize your instruments by heating them to 120°C. You rummage around in the kitchen cupboards and find a pressure cooker that can heat water to a pressure of 1.4 atm. Will this be sufficient for sterilizing your instruments? (You may use Table 5.6 in your book to answer this question; see the previous slide.) • Yes • No • Explain you answer.

23. 1 atm = 760 Torr Questions (Answer) • 1.4 atm (760Torr/atm) = 1064 Torr • This is less than the pressure required to reach 110°C (1075 Torr), therefore it is an insufficient pressure to reach 120°C. • (120-100)/(125-100)*(1741-760)+760=~1544 Torr (interpolation)

24. Solutions • Biological systems are mixtures of substances • Pure substances contain only one type of element or compound • They contain only one type of atom or molecule: • H2 • Hg • O2 • H2O • sucrose (C12H22O11) • Mixtures contain more than one type of pure substance • Heterogeneous mixture - components are not evenly mixed at the molecular level. • Homogeneous mixture - components are evenly mixed at the molecular level.

25. Solutions • A solution is another name for homogeneous mixture. • Solvent - the major component in a solution • Solute - the minor component in a solution.

26. Solutions • A solution is another name for homogeneous mixture. • Liquid solutions should be clear (transparent). • Liquid solution’s solutes should not settle with time • This distinguishes solutions from suspensions and colloids.

27. Solutions • In order form a solution to form • the solute molecules have to be able to form similar noncovalent interactions with the solute molecules as • the solute molecules form with themselves • the solvent molecules form with themselves.

28. States of Matter • Simulation of Glycerol and PropaneDissolving in Water

29. Solutions • Solubility is a measure of how much solute will dissolve in a solvent. • Solubility depends on temperature. • The solubility of gases decrease with increasing temperature • The solubility of solids and liquids usually increase with increasing temperature.

30. Solutions • When a solution is saturated, the solute dissolves and precipitates at the same rate.

31. Solutions • In order form a solution to form • the solute molecules have to be able to form similar noncovalents with the solute molecules as • the solute molecules form with themselves • the solute molecules form with themselves.

32. Solubility of Gases in Water • Henry’s Law - The solubility of a gas in a liquid is proportional to the pressure of the gas over the liquid. • The fizzing of soda when the cap is removed is an example of the lowered solubility of CO2 in water when it’s pressure above the soda is descrease. • The solubility of CO2 in water is very high, because it can react with water to produce and even more soluble product, H2CO3 (carbonic acid): • We will see that this is a very important reaction in biochemistry

33. Organic Compounds • Nonpolar, organic solutes will dissolve readily in nonpolar, organic solvents. • “Like dissolves Like”

34. Organic Compounds • The solubility is determined by the balance between the polar and nonpolar portions of the molecule.

35. Biochemical Compounds &Their Interactions with Water • Biological molecules are grouped into three categories. • Hydrophilic (water loving) molecules. • Polar molecules that can interact favorably with water • Hydrophobic (water fearing) molecules. • Nonpolar molecules that cannot interact favorably with water • Amphipathic molecules, which are conflicted about their feelings towards water. • Molecules containing both very polar and very nonpolar parts.

36. Biochemical Compounds &Their Interactions with Water • Hydrophilic (water loving) molecules. • Polar molecules that can interact favorably with water • Carbohydrates (sugars) have lots of polar hydroxyl groups

37. Biochemical Compounds &Their Interactions with Water • Hydrophilic (water loving) molecules. • Polar molecules that can interact favorably with water • Amino acids have both an amino and a carboxylic acid group, which are polar.

38. Biochemical Compounds &Their Interactions with Water • Hydrophobic (water fearing) molecules. • Nonpolar molecules that cannot interact favorably with water • The carboxylic acid groups, though polar, are dominated by the long hydrocarbon portions

39. Biochemical Compounds &Their Interactions with Water • Hydrophobic (water fearing) molecules. • A nonpolar solute "organizes" water • The H-bond network of water reorganizes to accommodate the nonpolar solute • This is an increase in "order" of water-This is a decrease in ENTROPY

40. Biochemical Compounds &Their Interactions with Water • Hydrophobic (water fearing) molecules.

41. Biochemical Compounds &Their Interactions with Water • Amphipathic molecules, which are conflicted about their feelings towards water. • Molecules containing both very polar and very nonpolar parts.

42. Biochemical Compounds &Their Interactions with Water • When placed in water, amphipathic molecules, form structures, such as micelles, which attempt to address the conflict.

43. Colloids and Suspensions • It is also possible to have mixtures which are not uniform at the molecular level. • These are called heterogeneous mixtures. • When a heterogenous mixture involves the mixing of a solid with a liquid, there are two possible situations: • Suspensions: • With time the solid settles out of the mixture • Colloids: • The solid stays suspended in the liquid indefinitely, • Both suspensions and colloids are cloudy.

44. Colloids and Suspensions

45. Diffusion and Osmosis • Within a solution, the solute and solvent molecules are constantly moving • If the concentration of the solute is not uniform throughout a solution, this movement will cause a net movement of solute molecules from the regions of high concentration to the regions of low concentration • In the end the concentration will be the same everywhere.

46. Diffusion and Osmosis • This movement is called diffusion.

47. Diffusion and Osmosis • If a semipermeable membrane that only allows solvent to pass through it is used to separate a region of high solute concentration from a region of low solute concentration • Solvent will move through the membrane from the region of low solute concentration to the region of high solute concentration in an effort to make the solute concentration the same on both sides of the membrane.

48. Diffusion and Osmosis • This movement is called osmosis.

49. Diffusion and Osmosis • This movement can be stopped by applying a pressure to the surface of the solution on the high solute concentration side of the membrane. • The pressure required to stop the movement is called the osmotic pressure.

50. Diffusion and Osmosis • Osmotic pressure is an important concept for understanding biological systems because the cell membrane is a semipermeable membrane • If the solute concentrations are not equal on both sides of the membrane, the cells can either shrivel up or swell up and explode