Name ____________________________________ Pd ____ Date ______________________ Changes of State Guided Inquiry Changes of State -­‐ There are 6 types of state changes: 1. Write a definition and give an example of each state change. a. Vaporization (aka boiling) b. Condensation c. Melting d. Freezing e. Sublimation f.
Deposition Energy, State Changes & Temperature State changes either require energy or release energy when they occur. When a particle goes from a lower energy state to higher energy state energy is required. Conversely, when a particle goes from a higher energy state to lower energy state energy is released. When energy is required during a state change the energy stored as interaction energy, Ei, is increased. When energy is released during a state change the energy stored as interaction energy is decreased. The absorption or release of energy must occur during a state change because of the conservation of energy law. 2. Which changes of state require energy? 3. Which changes of state release energy? The amount of energy required or released is the enthalpy of the state change: •
Energy required to change liquid to gas = enthalpy of vaporization = ΔHv •
Energy required to change solid to liquid = enthalpy of fusion = ΔHf •
Energy required to change solid to gas = enthalpy of fusion + enthalpy of vaporization = ΔHf + ΔHv 4. Look up the enthalpy of fusion and enthalpy of vaporization values on the Internet for the substances below. Include the proper units. Enthalpy of Fusion, ΔHf Enthalpy of Vaporization, ΔHv Neon Oxygen Water Increasing Energy → 5. Rank the state changes based on the amount of energy they require or release. Write a state change equation that shows the state change and quantifies the amount of energy required or released in terms of ΔHv and ΔHf. The melting and condensing state change equations are filled in as examples. Condensation Melting Solid + ΔHf → Liquid Gas → Liquid + ΔHv When energy is added to a particle it will either store the energy using the thermal energy storage mechanism, Eth, or the interaction energy storage mechanism, Ei. Increasing energy stored as thermal energy raises the temperature of the particle and increases its speed of motion. Increasing energy stored as interaction energy changes the particle to a higher state and increases its range of motion. The particle can switch between storage mechanisms to store incoming energy, but it can only use one storage mechanism at a time. In other words, a particle can’t store energy as thermal energy and interaction energy at the same time. 6. Reflect back on your Change of State Investigation and for each situation below indicate what happens to temperature (T ↑, T ↓ or T ↔), thermal energy (Eth ↑, Eth ↓ or Eth ↔), and interaction energy (Ei ↑, Ei ↓ or Ei ↔) when energy is added or removed from the particle. State(s) of Substance Heating (adding energy) Cooling (removing energy) Solid Solid + Liquid Liquid 7. Compare and contrast vaporization and condensation. 8. How do the boiling point and condensation point temperatures of a substance compare to each other? 9. Compare and contrast freezing and melting. 10. How do the freezing point and melting point temperatures of a substance compare to each other? 11. Compare and contrast sublimation and deposition. 12. How do the sublimation point and deposition point of a substance compare to each other? Energy, State Changes & Pressure We know from experience that pressure forces objects closer together. The kinetic molecular theory predicts that increasing pressure will hold particles closer together. Pressure acts to keep particles together just like the particles’ natural attractive forces keep the particles together. So as the pressure increases the forces holding the particles together get stronger. It takes more energy and a higher temperature to get the particles to change the solid state to the liquid state and from the liquid state to the solid state. Lowering the pressure has the reverse effect. It is easier for the particles to move farther apart, so less energy and a lower temperature is needed to go from the solid state to the liquid state and from the liquid state to the solid state. 13. Would you expect the boiling point of water in Denver, CO (elevation ~1,600 meters) to be the same, higher or lower than the boiling point of water in Annville, PA (elevation ~ 150 meters)? Why? 14. Would you expect the boiling point of water at the bottom of the ocean to be the same, higher or lower than the boiling point of water at the surface of the ocean? Why? Equilibrium Equilibrium is a dynamic condition in which two opposing changes occur at equal rates in a closed system. For example, a mixture of liquid water and ice is at equilibrium if the amounts of liquid water and ice do not change. When we look at a system in equilibrium it appears to us that the particles are not changing from one phase to the other. However, that is not the case. Particles are changing from solid to liquid and from liquid to solid at equilibrium. However, for every particle that changes from solid to liquid there is another particle that changes from liquid to solid. So there is no net change in the amount liquid water and ice. Show on Right: Equilibrium -­‐ Particles going from solid to liquid state at the same rate they are going from liquid to solid state. Both states will exist indefinitely. Two states (aka phases) can be present at the same time, but not be in equilibrium. This is called disequilibrium. For example, a mixture of liquid water and ice is in disequilibrium if either more liquid water is forming (the ice is melting) or if more ice is forming (the liquid water is freezing). In the case where the ice is melting it appears to us that only particles are changing from the solid phase to the liquid phase. However, particles are also changing from liquid to solid. Since particles are changing from solid to liquid faster than the particles are changing from liquid to solid, the net change is an increase in the amount of ice and a decrease in the amount of liquid water. The reverse is true when the water is freezing. Particles are still going from solid to liquid, but at a slower rate than the particles that are going from liquid to solid. Disequilibrium -­‐ Particles going from liquid to vapor state faster than they are going from vapor to liquid state. Eventually all the particles will be vapor. Disequilibrium -­‐ Particles going from vapor to liquid state faster than they are going from liquid to vapor state. Eventually all the particles will be liquid. 15. Water and ice cubes are kept in a glass at 0°C. a. Are the water particles melting faster, freezing faster or are both processes happening at the same rate? b. Are the two phases in equilibrium or disequilibrium? c. What will happen to the water and ice after a period of time? 16. Water and ice cubes are in a glass that is in a room with a temperature of 25°C. a. Are the water particles melting faster, freezing faster or are both processes happening at the same rate? b. Are the two phases in equilibrium or disequilibrium? c. What will happen to the water and ice after a period of time? 17. A pot of water is placed in an oven at 125°C. a. Are the water particles vaporizing faster, condensing faster or are both processes happening at the same rate? b. Are the two phases in equilibrium or disequilibrium? c. What will happen to the water after a period of time? Evaporation Evaporation is the process by which particles escape from the surface of a non-­‐boiling liquid and become gas particles. This process is slightly different than vaporization (boiling) because evaporation takes place at temperatures below the boiling point, whereas vaporization is the process whereby liquid particles become gas particles at the substances boiling point. The most common example of evaporation is when things “dry out”. If you spill water on the floor it will eventually dry, which means all the liquid particles entered the gas phase. This takes place even though the temperature never reaches the boiling point of water (100°C). How can a particle enter the gas phase if the temperature is below its boiling point? At a given temperature, a liquid has an average Eth and Ei, but some particles have more energy and some have less energy. The high energy particles have enough Ei to escape into the gas phase even though the temperature is below the boiling point. As the temperature increases the number of particles with enough Ei to enter the gas phase increases and so does the evaporation rate. 18. Where does evaporation occur in the liquid-­‐vapor system? 19. How can particles go from the liquid phase to the gas phase without being at the boiling point? 20. There are 3 different liquids, how can we use their boiling points to determine which one will evaporate the fastest? Vapor Pressure The pressure exerted by a vapor in equilibrium with its corresponding liquid is called the equilibrium vapor pressure of the liquid. The vapor particles exert a pressure that is proportional to their concentration in the gas phase. If some liquid is put in a container with a lid (a closed system) the liquid will begin to evaporate. Initially, particles leave the liquid state at a faster rate than the particles returning to the liquid state. The system is in disequilibrium and the vapor pressure continues to increase as more of the liquid evaporates. As more and more particles enter the gas state, the rate of particles leaving the gas state and returning to the liquid state increases. Eventually the rate at which particles leave and return to the liquid state are equal and the system is in equilibrium. The vapor pressure when the system is in equilibrium is called the equilibrium vapor pressure. Vapor pressure increases as temperature increases because more liquid particles have enough Ei to enter the gas phase. When the temperature reaches the normal boiling point the vapor pressure will equal the atmospheric pressure (760 torr). Under these conditions all the particles have enough Ei to enter the gas phase. For a given temperature, the lower the liquid’s boiling point the higher its vapor pressure. 21. At 20°C, what is the vapor pressure of diethyl ether, ethyl alcohol and water? 22. At 40°C, what is the vapor pressure of diethyl ether, ethyl alcohol and water? 23. What is the relationship between equilibrium vapor pressure and temperature? Temperature, Pressure & Phase Diagrams Phase diagrams are used to show the state of a substance for any given combination of temperature and pressure. Each substance has its own unique phase diagram. Phase Diagram for Water Phase Diagram for Carbon Dioxide The curves that separate the states on the phase diagram represent the combinations of temperature and pressure that create equilibrium between the states. The curve between the solid and liquid states represents the melting/freezing points of a substance. Notice the solid-­‐liquid curve for water has a negative slope and the solid-­‐
liquid curve for carbon dioxide has a positive slope. The solid-­‐liquid curve has a positive slope for most substances. Water is the exception! The negative slope for water means that liquid water is denser than ice and that increasing pressure lowers the melting point instead of raising the melting point. The curve between the liquid and vapor states represents the boiling/condensation points of a substance. Notice that increasing the pressure increases the temperature for a substance to boil/condense. The curve between the solid and vapor states represents the sublimation/deposition points of a substance. For water sublimation only happens at pressures less than 0.006 atmospheres; whereas with carbon dioxide sublimation happens at pressures up to 5.11 atmospheres. Since a substance’s boiling and melting/freezing points depends on the pressure, we use the terms normal melting/freezing point and normal boiling point to mean the temperature at 1 atmosphere of pressure. There are two other important points on a phase diagram. The triple point is the temperature and pressure where all three states of a substance are in equilibrium with each other. The critical point is the temperature and pressure above which the gas and liquid states become one state called a supercritical fluid. 24. What state is water in at 100 ℃ and 1.2 atmospheres? 25. What state is water in at 0 ℃ and 0.001 atmosphere? 26. What state is carbon dioxide in at 398.15 K and 60 atmospheres? 27. What state is carbon dioxide in at 190 K and 1.0 atmosphere? 28. Using the water phase diagram what states are in equilibrium at … a. point B? b. point F? c. point T? 29. What state is water in at 375 °C and 220 atmospheres? 30. What state is carbon dioxide in at 305 K and 75 atmospheres? Energy Calculations for State Changes We can use the enthalpies of vaporization (ΔHv) and fusion (ΔHf) to calculate the amount of energy that is required or released to change the state of a known mass of water; or we can calculate the mass of water that changed state if we know how much energy was absorbed or released. The state change energy equations are: for melting or freezing ∆𝐸! = 𝑄 = 𝑚×∆𝐻! for vaporization or condensation ∆𝐸! = 𝑄 = 𝑚×∆𝐻! for sublimation or deposition ∆𝐸! = 𝑄 = 𝑚× ∆𝐻! + ∆𝐻! For water at 1 atm ΔHf = 333.4 J/g ΔHv = 2266 J/g Sample Problem #1 – How much energy is required to completely vaporize 100.0 g of water? 𝑄 = 𝑚×∆𝐻! = 100.0 𝑔 × 2266 𝐽
= 226 400 𝐽 1 𝑔𝑟𝑎𝑚
Sample Problem #2 – It took 100.0 kJ of energy to melt a piece of ice. What was its mass? 𝑀𝑎𝑠𝑠 =
𝑄
1000 𝐽
1 𝑔
= 100.0 𝑘𝐽 𝑥 ×
= 299.9 𝑔 ∆𝐻!
1 𝑘𝐽
333.4 𝐽
31. How much energy is given off when 250.0 g of water is frozen? 32. How many grams of water can 479.0 kJ of energy vaporize? 33. How much energy is required to vaporize (boil) 75.0 g of water? Combined Energy Calculations When a substance is heated or cooled it may change temperature as well as state. We can use our combined energy equation to calculate the amount of energy required to change the temperature and state of a substance. 𝑄 = ∆𝐸!! + ∆𝐸! = 𝑚×𝐶×∆𝑇 + 𝑚×∆𝐻 temp change + state change The specific heat of a substance is different for different states. Be sure you are using the correct specific heat. For example, the specific heats for ice, water and water vapor are different. If the temperature change is going through several state changes we must expand the equation: 𝑄 = 𝑚×𝐶!"#$% ×∆𝑇!"#$% + 𝑚×∆𝐻! + 𝑚×𝐶!"#$"% ×∆𝑇!"#$"% + 𝑚×∆𝐻! + 𝑚×𝐶!"# ×∆𝑇!"# solid temp change solid-­‐liquid state change temp change liquid-­‐gas state change gas temp change Sample Problem How much energy is required or released to heat 50.0 g of ice from -­‐20.0°C to 120.0°C steam? Cice = 2.03 J/g•°C Cwater = 4.184 J/g•°C Csteam = 1.89 J/g•°C ΔHf = 333.4 J/g ΔHv = 2266 J/g Energy to heat ice + energy to melt ice + energy to heat water + energy to vaporize water + energy to heat steam Total Energy 50.0 g x 2.03 J/°C•g x (0°C – (-­‐20.0°C)) + 50.0 g x 333.4 J/g + 50.0 g x 4.184 J/°C•g x (100.0°C – 0°C) + 50.0 g x 2266 J/g + 50.0 g x 1.89 J/°C•g x (120.0°C – 100°C) = 2030 J = 16700 J = 20900 J = 113000 J = 1890 J 155000 J required 34. How much energy is required or released to cool 25.0 g of steam from 105°C to ice at -­‐5.0°C?