1. Entropy Review  Determine whether the entropy is increasing or decreasing for each of the following and why?  Loading the dishwasher  Condensation of water  Nuclear Fusion ( Inc / Dec )

2. Pressure  Pressure is the force per unit area exerted by a solid, liquid or gas  In KMT, pressure is described as being caused by the collision of molecules with their container.  Standard SI unit = Pascals (Pa)

3. Atmospheric Pressure  Atmospheric Pressure is considered to be a relative constant = 101,300 Pa  Atmospheric Pressure in various units:  Atmospheres → 1 atm  Kilopascals → 101.3 kPa  mm Hg or torr → 760 mm Hg or 760 torr  pounds per square inch → 14.7 psi  Convert the following:  1.2 atm into Pascals  769 torr into atm

4. Temperature  Temperature is related to the speed of the molecules  The standard SI unit for temperature is the Celsius (°C) or the Kelvin (K)  In calculations for gases, you must use the KELVIN scale!

5. Absolute Zero  At absolute zero (0 K), all molecular movement stops.  0 K = -273 °C

6. Temperature Conversions  Convert from Celsius to Kelvin:  °C = K – 273  K = °C + 273  Convert from Fahrenheit to Celsius:  (°C x 9/5) + 32 = °F  °C = (°F – 32) x 5/9  Convert the following:  32 °F to °C  53°C to °F  1513 C° to K  -44 °C to °F  146 K to °C  1200 F° to K

7. Temperature and Pressure  From the triple point graph, what happens to boiling point if  a) pressure decreases?  b) pressure increases?

8. Factors affecting reaction rates  A manometer measures experimental pressure whereas a barometer measures atmospheric pressure.

9. Vapor Pressure  Vapor Pressure is defined as the pressure exerted by a vapor in thermodynamic equilibrium with its condensed phases (solid or liquid) at a given temperature in a closed system.  Liquids that evaporate easily tend to have high vapor pressures.

10. Gas Laws  Gases have a very predictable relationship between their pressure, volume and temperature since the gas molecules are not bonded together.  The COMBINED GAS LAW relates these three variables in a closed gas system where no gas can enter or leave the system.  Units on both sides of the equation must match and you must use temperature in KELVIN.

11. Gas Laws  The volume of a gas-filled balloon is 30.0L at 40.0 °C and 153 kPa. What volume will the balloon have at STP?  STP Pressure = _______ Temperature = __________  A 2.45L football initially inflated to 1.25atm is brought in from the cold (4°C). What is the final volume when it is inside at room temperature?

12. Boyle’s Law  We often see the laws split into three separate laws: Boyle’s, Charles’ and Gay-Lussac  Boyle’s Law:  Constant temperature means isothermal  This is an Inversely Proportional relationship.

13. Boyle’s Law  The balloon has been filled with 3.0 L of air. Being creative, they try to “squish” it to half its original size. If the pressure begins at 18 psi, what will the final pressure be?  At room temperature, a 8.57L container begins with a pressure of 215kPa, What will the final volume be at 1.85atm?

14. Charles Law  Constant Pressure means isobaric  This is a directly proportional relationship

15. Examples:  A 1.75 L balloon at 25°C was placed into a freezer at -2°C. What will the final volume be?  2. What is the initial temperature of a 4.5L basketball when taken out of a cooler? The final temperature is 29°C and the volume increases to 5.1L

16. Gay-Lussac’s Law  Constant Volume means isovolumetric  The relationship is directly proportional

17. Ideal Gas Law:  Ideal Gas Law: The ideal gas law is the only law that allows the AMOUNT of gas to change.  The ideal environment for any gas is under LOW pressure and HIGH temperature.  P = Pressure in atm  V = Volume in liters  T = Temperature in Kelvin  n = moles in mol  R = 0.082 in L atm/mol K

18. Examples  Find the pressure of an ideal gas given a temperature of 18°C, 5.5g of chlorine, and the gas is 550 mL in size, and the R value is 0.0821 L atm/mol K.  3.8L of Nitrogen gas is at 163 kPa and 64°C. Assuming this is an ideal gas, how many grams of nitrogen are present?

19. Dalton’s Law of Partial Pressures  Partial Pressures: The law of partial pressures state that the sum of all partial pressures is equal to the total pressure of a system.  1. Three cylinders of equal volume are added together. The first cylinder is at 2.3atm, the second is at 150kPa, and the third is at 900mmHg. What is the total pressure?  2. The total pressure of a Nitrogen chamber is 3.8 atm. If 195kPa of nitrogen is extracted, what is the partial pressure of the nitrogen left in the chamber?

20. Effect of Concentration  Earlier we said that if we increase the concentration of a species, the reaction will try to reverse that increase.  PCl 5 (g) ↔ PCl 3 (g) + Cl 2 (g)  Adding PCl 5 will cause the system to produce less PCl 5 (→)  Removing PCl 5 will cause the system to produce more PCl 5 (←)

21. System in Equilibrium  The following system is at equilibrium. How would each of the following changes shift the position of the equilibrium?  CO (g) + 2H 2 (g) ↔ CH 3 OH (g) ∆H = -91 kJ  What happens if…..  Raise the temperature.  Increase the size of the container.  Increase pressure.  Add solid Fe mesh to absorb CO(g).  Add more H 2 (g).  Add more Carbon Monoxide

22. Equilibrium Constants Calculations ~ K eq  The Equilibrium Constant, K c or K eq, relates to a chemical reaction at equilibrium.  It can be calculated if the equilibrium concentration of each reactant and product in a reaction at equilibrium is known.

23. Equilibrium Constants Calculations ~ K eq  Below is a “generic” balanced chemical equation: aA + bB  cC + dD  Lower case letter represent the coefficients.  Upper case letters represent the molecules/atoms  The equation for equilibrium constant  Brackets mean concentration in MOLARITY

24. More Equilibrium Constant  Pure solids and liquids should be “excluded” from the equilibrium expression by placing a “1” If K >1 products are favored at equilibrium If K < 1 reactants are favored at equilibrium  A Homogenous equilibrium has everything present in the same phase. The usual examples include reactions where everything is a gas, or everything is present in the same solution.  A Heterogeneous equilibrium has things present in more than one phase. The usual examples include reactions involving solids and gases, or solids and liquids.

25. Examples  Gaseous Dinitrogen Tetroxide is in equilibrium with Gaseous Nitrogen Dioxide. Write and balance this equation, then write the equilibrium expression.

26. Example (cont’d)  Calculate K eq for the reaction in question 2 above when [SO 3 ] = 0.016M, [SO 2 ] = 0.0056M, and [O 2 ] = 0.0021M.  Does this reaction favor reactants or products?

27. Equilibrium-Pressure  You can also write an equation using the pressures of the reactants and products.  The pressure constant and the equilibrium constant are related. R = 0.0821 atm*liter / mol*K

28. Thermodynamics  Movement of heat energy  Specific Heat (or Specific Heat Capacity = C ) the amount of heat necessary to move 1.00 gram of a substance 1.00 °C  Heat Capacity is the amount of heat necessary to move the temperature 1.00 °C  DIFFERENT SOURCES DEFINE THESE DIFFERENTLY – I AM SORRY….SCIENCE IS WEIRD

29. Thermodynamics  Q = Heat Energy in Joules or calories  Specific Heat Capacity of Water  1 cal/g °C  4.184 J/g °C

30. Thermodynamics  1. A hot iron bar is thrust into 200-mL of water at 25 °C. If the water rises to 42 °C, what is the amount of heat energy gained by this endothermic reaction?

31. Thermodynamics  2. In an exothermic reaction, 50-kilocalories are emitted from a hot brick having a heat capacity of 1.32 cal/g °C. What was the mass of the brick if it began at 300 °C and fell to 45 °C?

33. Energy for Phase Changes  Heat of Fusion (H f ) or Molar Heat of Fusion is the energy absorbed to melt/or released to freeze per unit mass  Water H f = 6.01 kJ/mol  Water H f = 3.34 x 10 2 J/g  Heat of Vaporization (H v ) or Molar Heat of Vaporization is the energy absorbed to boil/or released to condense per unit mass  Water H v = 40.7kJ/mol  Water H v = 2.26 x 10 3 J/g

34. Energy for Phase Changes  If you start with 5.0g of ice at 0°C, how much energy must be absorbed to melt it all?  2. How many grams of ice can be melted with the addition of 1.5kJ of energy?  3. How many moles of water can be evaporated with the addition of 20kJ of energy?

35. Energy for Phase Changes

37. Phase Change Graph  Show on the graph the location of the boiling point temperature and the melting point temperature.  What phase of matter is missing? Name and describe it.  What do you note about temperature during the phase change?