2. Gas Relationships Gas Laws

3. Gas Variables  Temperature (K) = avg Kinetic Energy Kelvin = C + 273 Kelvin = C + 273  Volume = length x width x height (ml or cm 3 )  Pressure = force/Volume (atm, mmHg, or kPa)  Amount of Matter = number of moles

4. Kinetic Theory of Matter Gases consist of large numbers of molecules (or atoms, in the case of the noble gases) that are in continuous, random straight line motion  The volume of all the molecules of the gas is negligible compared to the total volume in which the gas is contained (This is equivalent to stating that the average distance separating the gas particles is large compared to their size).  Attractive and repulsive forces between gas molecules is negligible (they can never group together and form drops and liquefy)  Energy can be transferred between molecules during collisions, but the collisions are perfectly elastic and the total energy remains constant. Video Demo: Video Demo: http://comp.uark.edu/~jgeabana/mol_dyn/KinThI.html http://comp.uark.edu/~jgeabana/mol_dyn/KinThI.html http://comp.uark.edu/~jgeabana/mol_dyn/KinThI.html

5. Ideal Gas  Particles have no volume  Particles are in: Constant, rapid, random motion Constant, rapid, random motion Always move in straight lines Always move in straight lines  No attractive or repulsive forces  Temperature (K) proportional to Kinetic Energy

6. Real Gases  Besides the previous assumptions not being true:  Real Gases: Liquefy under High Pressure and Low Temperature.  Ideal Gases: Work best under Low Pressure and High Temperature.

7. Standard Temp and Press (STP)  273 K and 1 atm  273 K and 101.3 kPa  273 K and 760 mm Hg

8. Gas Laws  Boyle’s Law  Charles’ Law  Gay-Lusac Law  Dalton’s Law  Combined Law  Ideal Law  P1V1 = P2V2  V1/T1 = V2/T2  P1/T1 = P2/T2  Pt = P1 + P2 + ….  P1V1/T1 = P2V2/T2  PV = nRT

9. Pressure Versus Volume P 1 V 1 = P 2 V 2  Pressure Increases-Volume Decreases  Pressure Decreases-Volume Increases  http://www.kentchemistry.com/links/Gas Laws/Boyles.htm

10. Volume Versus Temperature V 1 /T 1 = V 2 /T 2  Volume Increases-Temperature Increases  Volume Decreases-Temperature Decreases  Temp in Kelvin  http://www.kentchemistry.com/links/Gas Laws/charles.htm http://www.kentchemistry.com/links/Gas Laws/charles.htm http://www.kentchemistry.com/links/Gas Laws/charles.htm

11. Pressure Vs Temperature P 1 /T 1 = P 2 /T 2  Pressure Increases-Temperature Increases  Pressure Decreases-Temperature Decreases  Temp in Kelvin  Kelvin = C + 273  http://www.kentchemistry.com/links/GasL aws/GayLussac.htm http://www.kentchemistry.com/links/GasL aws/GayLussac.htm http://www.kentchemistry.com/links/GasL aws/GayLussac.htm

12. What is the Paradox?  In  In looking at these Gas Laws a Paradox emerges:  As  As Pressure goes UP, Volume Goes DOWN Volume goes DOWN, Temperature goes DOWN Temperature goes DOWN, Pressure goes DOWN  How  How is that possible? Pressure went UP to start with?

13. Combined Gas Law P 1 V 1 /T 1 P 1 V 1 /T 1 = P 2 V 2 /T 2  Real World: You change one variable - ALL Change  Temp must be in Kelvin

14. Partial Pressures Pt = P1 + P2 + …..  Total Pressure = Adding up the Parts  http://demonstrations.wolfram.com/Dalt onsLawOfPartialPressures/ http://demonstrations.wolfram.com/Dalt onsLawOfPartialPressures/ http://demonstrations.wolfram.com/Dalt onsLawOfPartialPressures/  http://www.chm.davidson.edu/vce/GasL aws/DaltonsLaw.html http://www.chm.davidson.edu/vce/GasL aws/DaltonsLaw.html http://www.chm.davidson.edu/vce/GasL aws/DaltonsLaw.html

15. Avogadro’s Principle  Equal volumes of gases under the same conditions have:  Equal number of moles

16. Graham’s Law of Diffusion  http://www.youtube.com/watch?v=GRcZNCA9DxE http://www.youtube.com/watch?v=GRcZNCA9DxE  Lighter gases diffuse (travel) faster then heavier gases.  where: –r1 is the rate of effusion of the first gas (volume or number of moles per unit time). –r2 is the rate of effusion for the second gas. –M1 is the molar mass of gas 1 molar massmolar mass –M2 is the molar mass of gas 2.

17. Intermolecular Forces What binds molecules together?

18. INTRAMOLECULAR VS INTERMOLECULAR  Intramolecular force - Atoms within a molecule are attracted to one another by the sharing of electrons.  Intermolecular forces – Forces that hold molecules together in solids and liquids. These are much weaker then Intramolecular forces.

19. EXAMPLES OF INTRAMOLECULAR FORCES  Ionic Bonds – The attraction between ions to form compounds based on charge.  Covalent Bonds – The attraction between the nucleus of one atom and the electrons of another atom. The electrons are shared between atoms.

20. MORE ON COVALENT BONDS  Polar Covalent Bonds – The electrons in a covalent bond are shared unequally, giving one atom a slightly positive charge and the other a slightly negative charge.  Nonpolar Covalent Bonds – The electrons in a covalent bond are shared equally, resulting in both atoms having a neutral charge.

21. The States of Matter and Intermolecular Forces

22. The Nature of Intermolecular Forces  A very strong type of force that is responsible for much of chemistry is electrostatics.  The attraction of a positive charge with a negative charge is the force that allows for the structure of the atom,  causes atoms to stick together to form molecules, both ionic and covalent,  and is responsible for the formation of liquids, solids and solutions.

23. HOW INTERMOLECULAR FORCES WORK

24. Trends in the Forces  The intramolecular forces keep the atoms in a molecule together and are the basis for the chemical properties,  The intermolecular forces are those that keep the molecules themselves together and are virtually responsible for all the physical properties of a material.

25. INTRAMOLECULAR VERSUS INTERMOLECULER FORCES Covalent Bonds > Hydrogen Bonding> Dipole-Dipole Attractions> London Forces 400 kcal > 12-16 kcal > 2-0.5 kcal > Less than 1 kcal Clearly normal covalent bonds are almost 40 times the strength of hydrogen bonds. Covalent bonds are almost 200times the strength of dipole-dipole forces, and more than 400 times the size of London dispersion forces.

26. London Dispersion Forces  London dispersion forces exist in nonpolar molecules.  These forces result from temporary charge imbalances. The temporary charges exist because the electrons in a molecule or ion move randomly in the structure.  The nucleus of one atom attracts electrons from the neighboring atom.  At the same time, the electrons in one particle repel the electrons in the neighbor and create a short lived charge imbalance.

27. EXAMPLE These temporary charges in one molecule or atom attract opposite charges in nearby molecules or atoms. A local slight positive charge d+ in one molecule will be attracted to a temporary slight d- negative charge in a neighboring molecule.

28. Dipole-Dipole Attractions  Dipole-dipole attractions exist between molecules that are polar.  This requires the presence of polar bonds and an unsymmetrical molecule.  The slightly positive atom in one molecule is attracted to the slightly negative atom in another molecule.

29. EXAMPLE

30. ANOTHER EXAMPLE

31. Hydrogen Bonding  Hydrogen bonding is a unique type of intermolecular attraction.  There are two requirements: –The first is a covalent bond between an H atom and either F, O, or N these are the three most electronegative elements. –The second is an interaction of the H atom in this kind of polar bond with a lone pair of electrons on a nearby atom like F, O, or N.

32. WATER AND HYDROGEN BONDING  The normal boiling point for water is 100ºC.  The observed boiling point is high compared to the expected value.  The predicted boiling point from the trend of boiling points for H2Te, H2Se, H2S and H2O is very low.  If the trend continued the predicted boiling point would be below -62ºC.  The “anomalous” boiling point for water is the result of hydrogen bonding between water molecules.

33. WATER AND FREEZING  Hydrogen bonding is responsible for the expansion of water when it freezes.  The water molecules in the solid have tetrahedral arrangement for the two lone pairs and two single bonds radiating out from the oxygen.  The lone pairs on the “O” atom can attracted to nearby water molecules through hydrogen bonds.  A cage like structure results.

34. EXAMPLE OF ICE

35. Energy in Chemical Reactions

36. Potential Energy: Potential Energy is stored energy. It is energy something has because of its position or composition. What has more potential energy – a bowling ball on the ground, or a bowling ball three feet directly above your head? What has more potential energy – a barrel of water or a barrel of gasoline?

37. Chemical potential energy  That’s the type of energy that gasoline has.  Defined as the potential energy stored in the bonds between atoms of a substance.

38. Chemical potential energy  When you burn gasoline, or wood, or anything flammable, much of the chemical potential energy is released as heat.

39. Chemical potential energy  Burning isn’t the only chemical reaction that releases heat.  Virtually every chemical reaction either releases heat (exothermic) or absorbs heat (endothermic). Exothermic: heat is exiting.

40. Exothermic process is any process that gives off heat – transfers thermal energy from the system to the surroundings. Endothermic process is any process in which heat has to be supplied to the system from the surroundings. System? Surroundings? The “system” is generally the container where the reaction is taking place.

41. Exothermic process is any process that gives off heat – transfers thermal energy from the system to the surroundings. If energy is listed as a product, the reaction is exothermic: Endothermic process is any process in which heat has to be supplied to the system from the surroundings. If energy is listed as a reactant, the process is endothermic: 2H 2 ( g ) + O 2 ( g ) 2H 2 O ( l ) + energy energy + 2HgO ( s ) 2Hg ( l ) + O 2 ( g ) 6.2

42. Endothermic or Exothermic? (how to decide) Fires give off heat, and feel hot to your hands. They are exothermic. Cold packs feel cold to your hands, so they must be the opposite: endothermic. They feel cold because they’re absorbing heat from your hands.

43. Endothermic reactions feel cool because they absorb heat from the surroundings.

44. Chemical energy Is this reaction exothermic or endothermic?

45. Potential energy Heat Is this reaction exothermic or endothermic?

46. Chemical energy What about this one?

47. Potential energy Heat What about this one?

48. Practice Questions Is this reaction endothermic or exothermic? What is the change in energy for the reaction?

49. Practice Questions Does the graph represent an endothermic or exothermic reaction? What is the change in energy for the reaction?

50. Energy of chemical reactions: Every energy measurement has three parts: 1. A unit ( Joules or calories). 2. A number (how many). 3. and a sign to tell direction: negative = exothermic positive = endothermic

51. Enthalpy  The heat lost or gained by a system during a physical or chemical change is called the enthalpy change (ΔH) or the heat of reaction. Remember that a negative enthalpy change means that heat is lost from the system to the surroundings, making the process exothermic.  A positive enthalpy change tells us that heat is gained by the system and the process is endothermic.

52. System Surroundings Energy  <0

53. System Surroundings Energy  >0

54. Thermochemical Equations H 2 O ( s ) H 2 O ( l )  H = 6.01 kJ Is  H negative or positive? System absorbs heat; the system will feel cold. Endothermic  H > 0 6.01 kJ are absorbed for every 1 mole of ice that melts at 0 0 C and 1 atm. 6.4

55. Thermochemical Equations CH 4 ( g ) + 2O 2 ( g) CO 2 ( g) + 2H 2 O ( l )  H = -890.4 kJ Is  H negative or positive? System gives off heat Exothermic  H < 0 890.4 kJ are released for every 1 mole of methane that is combusted at 25 0 C and 1 atm. 6.4

56. H 2 O ( s ) H 2 O ( l )  H = 6.01 kJ The coefficients always refer to the number of moles of a substance Thermochemical Equations If you reverse a reaction, the sign of  H changes H 2 O ( l ) H 2 O ( s )  H = - 6.01 kJ If you multiply both sides of the equation by a factor n, then  H must change by the same factor n. 2H 2 O ( s ) 2H 2 O ( l )  H = 2 x 6.01 = 12.0 kJ 6.3

57. Practice Question Which statement correctly describes an endothermic chemical reaction? (1) The products have higher potential energy than the reactants, and the ΔH is negative. (1) The products have higher potential energy than the reactants, and the ΔH is negative. (2) The products have higher potential energy than the reactants, and the ΔH is positive. (2) The products have higher potential energy than the reactants, and the ΔH is positive. (3) The products have lower potential energy than the reactants, and the ΔH is negative. (3) The products have lower potential energy than the reactants, and the ΔH is negative. (4) The products have lower potential energy than the reactants, and the ΔH is positive. (4) The products have lower potential energy than the reactants, and the ΔH is positive.

58. Practice Question  What is  for this reaction? Number and sign)